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COPYRIGHT DEPOSIT. 



LECTURE NOTES 



ON 



GENERAL CHEMISTRY 



REVISED AND ENLARGED 



BY 



J. F. McGREGORY, A. M., F. C. S., 

PROFESSOR OF CHEMISTRY AND MINERALOGY IN 
COLGATE UNIVERSITY 



HAMILTON, N. Y.: 

REPUBLICAN PRESS. 

1902 



THE LIBRARY OF 
CONGRESS, 

Two Cowe8 Received 

JUN. U 1902 

E9HT ENTRY 
11 -,ctcr 
Sy)CXaNo. 

i ft V ^ 

COPY B. 



COPYRIGHT, 1894, 1902, BY 
J. F. McGREGORY. 



• ••• •••• •• • ••• ••• 



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PREFACE 



The first edition of these " Lecture Notes,'' published 
in 1894, having been exhausted, the author has revised 
and re-written them, and has added the notes on the 
metals contained in Part three, more than doubling the 
size of the book. 

It is the plan in operation in very many of our colleges 
to give instruction in Chemistry by experimental lectures 
accompanied by frequent examinations, and these 
"Notes" are designed to accompany such a course of 
lectures. The majority of college students have had so 
little practice in note taking, that they are quite likely 
to omit much that is important, while taking down that 
which is of less importance. These notes are intended to 
include the important facts which it is necessary for a 
student to learn, leaving him only the necessity of taking 
notes on the experiments and explanations given by the 
lecturer. It is the author's own plan to accompany the 
study of Part three with a laboratory course in element- 
ary Qualitative Analysis, covering the work as given in 
Part I of his " Manual of Qualitative Analysis." 

The author desires to express his thanks to his assist- 
ant, Mr. R. B. Smith, for his suggestions and help in the 
work of preparation, and to his colleague, Prof. R. W. 
Thomas for his careful reading and criticism of the man- 
uscript. 

J. F. M. 
Hamilton, N. Y., June, 1902. 



LECTURE NOTES 



-ON- 



GENERAL CHEMISTRY 



PART ONE 



Introduction 

1. Matter. All the objects by which we are sur- 
rounded, whatever else they may lack or however they 
may differ, possess the common property which we call 
weight. By this we mean that they are attracted, or 
drawn toward the earth, by something which we call 
gravitation. Weight is common not only to all solid sub- 
stances, such as stone, and to all liquids, such as water, 
but also to all gases, such as air. We use the word matter 
to designate all such bodies. Matter may therefore be 
denned as anything which has weight. It exists in three 
forms, the solid, the liquid and the gaseous. 

2. Force. Every form of matter is subject to changes 
which affect its condition. These changes may be either 
temporary or permanent, and can be shown to be directly 
connected with some movement in the particles of matter. 
That which imparts this motion to the particles and thus 



Z LECTURE NOTES 

causes the change, we call force. Force may also be em- 
ployed in changing or stopping motion in the particles. 
Force may therefore be defined as that which is capable of 
producing motion in the particles of matter, or of changing 
the direction or velocity of the particles of matter already 
in motion. 

3. Departments of Science. Matter may be studied 
in a great variety of ways. The knowledge of all the 
phenomena of the external world thus derived and sys- 
tematized, is classified under the general name of natural 

science. 

This knowledge may be gained by the study of external 
form or structure, of internal properties, of composition, 
or of the changes which take place when matter is acted 
upon by different forces. Many of the facts may be 
obtained by simple observation, while others require some 
test or experiment to prove them. 

Among the many departments of natural science we 
find two, Physics and Chemistry, which are so closely 
related as to form a group by themselves. They have to 
do with the external and internal properties of matter, 
and with the changes which matter undergoes when acted 
upon by forces. These are often called the physical 
sciences. 

4. Physical Changes. If matter is acted upon by 
forces so as to produce changes in condition, if, when the 
cause is removed, the substance resumes its original 
condition, the change, which is only temporary, is called 
a physical change. A change in the form of matter, such 
as the melting of ice or the boiling of water, is a physical 
change. The study of such changes, together with the 



LECTURE NOTES 6 

study of the forces which produce them, belongs to the 
science of physics. We may therefore define physics as 
that branch of physical science, which treats of the tempo- 
rary changes which take place when matter is acted upon 
by forces, and of the nature and properties of the forces 
which produce the changes. 

5. Chemical Changes. If matter is acted upon by 
forces so as to produce changes in condition, if, when the 
cause is removed, the substance remains permanently 
changed in condition, it is called a chemical change. 

A chemical change is always accompanied by a change 
in the weight of the substance acted upon. If the weight 
has increased, there has been a combination with some 
other substance. If the weight has decreased, a portion 
of .the substance has been liberated in the form of gas. 
There has been neither creation nor destruction of matter. 
This principle, which is universally true, is known as the 
indestructibility of matter. 

The study of such changes, as well as a knowledge of 
the nature of the substances which produce such changes, 
belongs to the science of chemistry. We may therefore 
define chemistry as that branch of physical science which 
treats of the ultimate composition of 'matter, of the perma- 
nent changes which take place when matter is acted upon by 
forces, and of the action and reaction of different kinds of 
matter on each other. 

6. Divisions of Matter. Modern scientists regard all 
forms of matter as made up of extremely minute particles. 
While we necessarily have to study matter in its aggre- 
gate form, which we call the mass, the belief is universal 
that these minute particles do exist, and indirect proof of 
their existence is not altogether lacking. These minute 



4 LECTURE NOTES 

particles are called molecules. Most molecules, under 
certain conditions, possess properties which would seem 
to indicate that they are composed of still more minute 
particles. As there is no indication that these latter 
particles can be further divided, they are called atoms. 
There are, therefore, three divisions of matter, the 
mass, the molecule and the atom. 

This view of the constitution of matter, together with 
certain laws which govern chemical combination, form 
what is known as the atomic theory, which is the founda- 
tion of all physical science. 

7. Mass and its Attractions. The mass is the only 
division of matter which can be appreciated by the senses. 
Masses differ very greatly in size, since the largest body 
in the universe, and the minutest portion that can 'tee 
appreciated by the most delicate instrument, each belongs 
to this division. 

All masses possess a certain amount of attraction for 
each other. In comparatively small bodies this is inap- 
preciable, but between the sun and the planets, as well as 
between smaller bodies and the earth, the amount is very 
great. This force is called gravitation, and it is this which 
gives matter weight. 

The measure of this force may be accomplished in two 
ways. We may compare the mass with a unit weight of 
some substance, which has been determined under fixed 
conditions, or, we may compare it with an equal volume 
of some standard substance under fixed conditions. The 
former is called absolute weight and the latter relative 
weight. 

8. Absolute Weights and Measures. There are many 



LECTURE NOTES O 

systems of weights and measures. The best and most 
extensively used is the French or metric system. It is a 
decimal system, which simplifies calculations, and is almost 
universally used in scientific work. 

In this system the standard unit is the mete?-, which is 
the unit of length, and from this all the other units in the 
system are derived. The meter is one ten millionth of a 
quadrant of a terrestrial meridian, as nearly as can be 
determined. The actual standard meter is the length of 
a bar of platinum, which is deposited in the Palace of the 
Archives of France in Paris. It is equal to 39.37 inches. 

The multiples and fractions of each unit in the system 
are denoted by prefixes placed before the names of the 
units. The multiple prefixes are derived from the G-reek, 
and are deka, ten, hecto, hundred, kilo, thousand, myria, 
ten thousand. The fractional prefixes are derived from the 
Latin, and are deci, tenth, centi, hundredth, and milli, 
thousandth. 

The unit of weight is the gram, which is equal to the 
weight of one cubic centimeter of pure water at 4° centi- 
grade under a barometric pressure of 760 millimeters. It 
is equal to 15. 43 grains. One ounce avoirdupois is equal 
to 28.35 grams, and one ounce Troy is equal to 31.1 
grams. 

The unit of capacity is the liter. This is equal to one 
cubic decimeter, or, as it is more commonly given, to one 
thousand cubic centimeters. It is equal to 1.056 U. S. 
standard liquid quarts, or to 61.02 cubic inches. 

9. Specific Gravity and Density. These are the terms 
which are used to express relative weight, that is, the 
weight of any substance compared with the weight of an 



b LECTURE NOTES 

equal volume of some substance which is taken as a 
standard. 

Specific gravity is the term used to express the relative 
weight of a solid or liquid, when compared with the weight 
of an equal volume of water ; or of a gas, when compared 
with the weight of an equal volume of air. 

Density is a term used to express the relative weight 
of substances in the gaseous state only. It is the relative 
weight of a gas, when compared with the weight of an 
equal volume of hydrogen. Hydrogen is used as the 
standard for density because it is the lightest substance 
known. 

The relation between specific gravity and density is 
dependent upon the relative weights of the two standards. 
By experiment we find that the density of air is 14.39. 
If, therefore, we multiply the specific gravity of a gas by 
14.39, we obtain its density ; or by dividing its density 
by 14.39, we obtain its specific gravity. 

10. The Effect of Heat on Matter. Heat causes all 
kinds of matter to expand. In the case of solids and 
liquids, the amount of expansion is very small and varies 
with the substance ; but in the case of gases the amount 
of expansion is much greater and the rate is found to be 
uniform. //' the temperature of the gas is reduced to 0° 
centigrade, the amount of expansion is ^i of the volume 
of the gas for each degree of increase in its temperature. 
This is known as Dalton's law. It is also often called the 
law of Charles, and of Gay-Lussac. 

The fraction, ^1-^, or its corresponding decimal, 0. 003665, 
is called the co-efficient of expansion for gases. 

Knowing this uniform rate of expansion, if we know 



LECTURE NOTES 7 

the volume which a gas occupies at a given temperature, 
we can find what the volume will be at any other temper- 
ature. For a unit volume at 0° will become 1 -4- pr^-r 

273 

at 1°, 1 + A at 2°, 1 + JL at t° and 1 + J*L at t x °. 

Let V represent the given volume at t°, and V x ]the re- 
quired volume at t x °. We shall then have the pro- 
portion, 

V:V X :: 1 + ^:1 + ^, or 

V : V x :: 273 -f t : 273 + t x . 
From this we obtain, 

/273 + t, 

\273 + t 

By substituting in this equation we obtain the required 
volume. 

11. Thermometers. A thermometer is an instrument 
for measuring temperature. In the construction of this 
instrument, the fact that liquids expand when heated, is 
the principle utilized. It consists of a very narrow glass 
tube with a bulb at one end, which is filled with some 
liquid, usually with mercury. The graduation of the tube 
is always made with reference to two fixed points, viz. : 
the melting point of ice and the boiling point of water. 

There are three important thermometric scales. These 
are, in the order of their appearing, the Fahrenheit, the 
Reaumur [ra-o-mur] and the Celsius, called also the 
Centigrade. 

In the Fahrenheit scale, the zero is the point at which 
the column of mercury stands when placed in a mixture 



8 LECTURE NOTES 

of common salt and ice. (1: 3). It was intended to have 
the normal temperature of the human body correspond to 
100° of the scale, and so the first graduation from these 
points, placed the melting point of ice at 32°, and the boil- 
ing point of water at 212°. After this thermometer was 
constructed, it was found that exact proportions in the 
mixture of salt and ice must be used, or the temperature 
would vary, and that the normal temperature of the human 
body was not 100° of this scale, but varied slightly, being 
98.4° to 98.6°. The two fixed points mentioned above are 
therefore always employed, the melting point of ice, (com- 
monly called the "freezing point") being marked 32°, and 
the boiling point of water, (commonly called simply the 
"boiling point,") marked 212°. This is the thermometer 
in common use in England and the United States. 

In the Reaumur scale, the freezing point is fixed at 0° 
and the boiling point at 80°. It is the thermometer in 
common use in Germany. 

In the Celsius, or Centigrade scale, the freezing point 
is fixed at 0° and the boiling point at 100°. It is the 
thermometer in common use in France. It is also uni- 
versally employed in scientific work and, being based on 
the decimal system, is the most rational and convenient 
scale in use. 

The relation of value between the degrees of these differ- 
ent scales, is the ratio of the number of degrees between 
the freezing and boiling points, which is that of 9 to 5 to 
4. [180°: 100°: 80°.] 

As it is often necessary to find the temperature in one 
scale from that in another, we may do so by means of the 
following equations, obtained from the above relations : 



LECTURE NOTES 9 

C o = 5 (F°— 32°). C° = f R°. 

F° =|C°| 32°. F° = fE°| 32°. 

R° = | (F° — 32°). R° = | C°. 

The degrees below the zero point are written with the 
minus sign preceding them. 

Absolute temperature is a term applied to temperatures 
reckoned from absolute zero, a temperature at which it is 
believed all motion and all chemical action ceases. This 
temperature is calculated to be — 273° Centigrade. 

12. The Barometer. We have already noted the fact 
that all forms of matter, including gases, have . weight. 
The atmosphere about us, therefore, has weight and must 
exert a certain amount of pressure upon the surface of 
the earth. Moreover, since the particles (molecules) of 
a gas move freely in all directions, atmospheric disturb- 
ances, as well as different elevations, must result in dif- 
ferent atmospheric pressures. 

The barometer is an instrument for measuring the 
atmospheric pressure. As ordinarily constructed, it con- 
sists of a glass tube, somewhat more than 800 millimeters 
long, closed at one end, filled with mercury and then 
inverted, the open end of the tube dipping into a cup of 
mercury. When this is done, the mercurial column, 
instead of filling the tube, will be seen to sink to a certain 
level. 

Since the atmosphere must press upon the surface of 
the mercury in the cup, and since the space above the 
column of mercury in the tube must be a vacuum, no air 
having been allowed to get into it, and since the particles 
of the liquid mercury must be in equilibrium, the weight 
of the mercury in the tube must just equal the weight of 



10 LECTURE NOTES 

the atmosphere on the same amount of surface. It is, 
therefore, a measure of the atmospheric pressure. 

Under ordinary conditions, and at the sea level, the 
column of mercury is 760 millimeters in height. This is 
called the pressure of one atmosphere, high pressure being 
often measured in atmospheres. If the interior of the 
glass tube has a cross-section equal to one square centi- 
meter, the weight of the mercury in the tube will exactly 
equal the atmospheric pressure on one square centimeter. 
The atmospheric pressure is found to be equal to 1033.3 
grams on one square centimeter, and this is equal to about 
15 pounds, avoirdupois, on one square inch. (14.7 pounds 
exactly.) This is the amount of one atmosphere of 
pressure. 

13. The Effect of Pressure on Matter. The effect of 
pressure on matter is to decrease its volume. In the case 
of solids and liquids the amount of this decrease is exceed- 
ingly small and varies with the substance. 

In the case of gases it is found that the volume is 
greatly decreased under pressure, that the decrease in 
volume is regular under a regular increase in pressure, 
and that the gases expand to their original volume when 
the pressure is removed. In so far as it is dependent 
upon pressure, the law of gases is, that the volume occupied 
by a given 'weight of any gas, is inversely prop>ortional to 
the pressure exerted upon it. This is known as Boyle's law. 

In order to change the volume of a gas from any given, 
to any required pressure, if we let V represent the given 
volume under the barometric pressure B, and V x the re- 
quired volume under the pressure B x , we have, according 
to Boyle's law: 

V B 
V : V x :: B x : B;or,V x = ■ ~ . 

B x 



LECTtJRE NOTES 11 

If we wish to combine this correction of pressure with 
that for temperature, we have the following equation : — 

f 273 4- .U. ) B 



273 + t ) B 



14. Molecules. Molecules are the minute particles 
which form the mass. Direct proof of their existence is 
not possible, nor is it necessary. There is hardly a 
chemical or physical phenomenon, that is not more clearly 
understood by knowing what the scientist means by the 
term molecule. 

If we heat a piece of white sugar, it soon becomes dark- 
colored, and in time, black. This black residue is carbon, 
which shows us that whatever else sugar may contain, 
one of its ingredients is carbon. If we grind a piece of 
sugar to powder, the finest particle will show the same 
result upon heating. We can easily understand that if 
it were possible for us to keep on sub-dividing the piece 
of sugar long enough, we should finally reach a particle, 
which, upon further subdivision, would give us carbon, or 
some other substance contained in the sugar, but not sugar 
itself. The ultimate particle of sugar we call the mole- 
cule. A molecule is, therefore, the smallest particle of 
matter which is capable of an independent existence. It is 
the smallest particle iixto which the mass can be divided 
without losing its identity. 

The actual size of molecules has been the subject of 
much speculation, but is a matter which cannot be deter- 
mined with exactness. Experiments by Lord Kelvin seem 
to show, that the distance between two adjoining molecules 
in a liquid, is somewhere between one ten-millionth and 
two hundred-millionths of a millimeter. In order to form 



12 LECTURE NOTES 

some conception of these minute particles, he calculated 
that if a drop of water should be magnified to the size of 
the earth, the molecules would be larger than shot but 
smaller than cricket-balls. 

Each molecule is composed of one or more atoms. If 
the molecule is composed of one kind of atoms only, it is 
called elementary ; if of more than one kind, it is called 
compound. 

Most elementary molecules consist of two atoms. 

15. Molecular Attraction. By molecular attraction 
we mean the force by which the molecules are held together 
in the mass. This force is entirely different from gravi- 
tation, which is the attractive force that holds masses 
together. Molecular attraction acts only through very 
minute distances, but within the sphere of its influence it 
is very strong. 

This force is called cohesion, when the molecules are of 
the same kind, and adhesion, when they are of different 
kinds. 

The amount of attraction which the molecules have for 
each other, differs with the kind of molecule but is strong- 
est in the metals. It also varies greatly with the form 
of matter in which the mass exists. 

In solids, the molecules are very near together, and 
molecular attraction is very strong. 

In liquids, the molecules are not so near together and 
consequently the molecular attraction is only very slight. 

In gases, the molecules are so far apart that molecular 
attraction has ceased. 

Heat tends to weaken the attraction between molecules, 



LECTURE NOTES 13 

and so changes solids to liquids and liquids to gases. 
These forms of matter are often called the states of ag- 
gregation of matter. 

16. Molecular Motion. At the ordinary temperature 
the molecules are in constant motion, whatever may be 
the state of aggregation of the mass. This motion in- 
creases as the temperature increases and, on the other 
hand, it is believed that all motion ceases at the absolute 
zero of temperature. [ — 273° C] 

The molecular motion in gases and liquids can be easily 
shown. That there is molecular motion in solids can be 
demonstrated in many cases, although in most solids the 
motion is too slow to make practical demonstration 
possible. 

Heat weakens molecular attraction and increases mole- 
cular motion. 

When a solid is heated, the molecules are driven farther 
and farther apart as the temperature increases. When 
the molecular attraction has been so weakened as to allow 
the molecules to move about freely among themselves, 
(though still held together by the attractive force of the 
molecules), we have a liquid. 

When a liquid is heated, the molecular motion becomes 
more and more rapid, and the molecular attraction becomes 
weaker and still weaker. When the kinetic energy of the 
moving molecules is sufficient to overcome entirely the 
molecular attraction, we have a gas. 

When a gas is heated, since the molecular attraction 
has now ceased, the only change that can occur is an 
increase in the rapidity of the molecular motion. If the 
gas is in a confined space this manifests itself in pressure 



14 LECTURE NOTES 

on the sides of the containing vessel. The heat necessary 
to produce these changes in matter varies greatly with 
the substance. 

The rapidity with which the molecules of a gas move 
is almost incredible. Maxwell estimates that the mole- 
cules of hydrogen at 0° move at the rate of about 1700 
meters per second. [1609 meters = 1 mile.] 

17. Diffusion of Liquids and Gases. Solid substances 
can only be mixed together by mechanical force. When 
two or more liquids are brought together, some mix 
readily, while others do not mix at all. If two liquids 
which mix are brought together carefully, they will 
remain separated for some time but will gradually become 
homogeneous. This gradual mixing is called diffusion. 

When gases are brought together, they always mix, 
and that quite rapidly, the rapidity being dependent upon 
their densities. This matter was investigated by Graham 
who found that gases diffuse themselves with a velocity 
which is inversely proportional to the square roots of their 
densities. 

This is known as Graham's law. 

18. Molecular Volume. By molecular volume we mean 
the amount of space which a molecule will occupy. fC This 
has nothing to do with the actual size of the molecule but 
depends entirely upon the state of aggregation of the 
mass. 

In solids and liquids, there is no general law of relation. 
The molecular volume is smaller in solids than in liquids, 
and in each it is smaller than in gases. 

The many ways in which gases are similarly affected, 
such as the uniform change in volume due to change in 



LECTURE NOTES 15 

temperature and pressure, leads to the very important 
deduction, that all molecules, in the gaseous state, occupy 
the same amount of space. 

This was first proposed as an hypothesis in 1811, by 
the Italian physicist, Avogadro, as the direct result of his 
study of the law of gaseous volumes. It has since been 
universally accepted as a fundamental principle of science 
and is known as Avogadro's law. It may be stated as 
follows : — Equal volumes of all kinds of matter in the gas- 
eous state, under the same conditions of temperature and pres- 
sure, contain the same number of molectdes. 

19. Molecular Weight. Although the molecule is so 
small that we can only approximate its absolute weight, 
we can obtain its relative weight with great exactness. 

We define the molecular weight of a substance as the 
weight of its molecule compared with the weight of an 
atom of hydrogen. Since hydrogen is taken as the stand- 
ard, the weight of its atom is one. We may also define 
molecular weight as the sum of the weights of the atoms 
which make up the molecule. Hydrogen is an elementary 
substance and its molecule is therefore composed of two 
atoms. This cannot, of course, be demonstrated, but all 
investigation tends to confirm the truth of the statement. 

If we wish to determine the molecular weight of a sub- 
stance in the gaseous state, we compare the weight of a 
certain volume of the gas with that of an equal volume 
of hydrogen, under the same conditions of temperature 
and pressure. This gives us the density of the gas. 
According to Avogadro's law, these equal volumes of 
gases, being taken under the same conditions of temper- 
ature and pressure, will contain the same number of 



16 LECTURE NOTES 

molecules. The density of a gas will therefore represent 
the weight of a molecule of the gas compared with the 
weight of a molecule of hydrogen. Since the hydrogen 
molecule, being composed of two atoms, weighs two, the 
molecular weight of any substance can be obtained by mul- 
tiplying its density by two. 

This method of obtaining the molecular weight can only 
be used for gases, or for such substances as can be changed 
into the gaseous state without decomposition. Chemical 
analysis gives us practical methods for determining mol- 
ecular weight which can be used in all cases, and there 
are several methods based upon the properties of substances 
when in solution. 

20. Latent Heat. We have already seen that heat 
expands matter. This it does by weakening the molecular 
attraction, and thus driving the molecules farther apart. 

The molecules, which are in a state of constant motion, 
can be acted upon by heat in two ways : — The velocity of 
their motion may be increased, the limits of their molecu- 
lar volume remaining the same ; or, the limits of their 
molecular volume may be increased, the velocity of their 

motion remaining the same. The former action results 
in an increase in the temperature of the body, and is called 
sensible heat, while the latter is always associated with a 
change in the state of aggregation, and, since it is not 
accompanied by a rise in temperature, it is called latent 
heat. Thus, a piece of ice at — 10°, subjected to heat, 
will grow warmer until it reaches 0°. If the heating be 
continued there will be no further change in temperature 
until all of the ice is melted. [Increased limit of mole- 
cular volume. J The heat which melts the ice is rendered 
latent. 



LECTURE NOTES 17 

If the water, formed, be now heated, the temperature 
will rise until it reaches 100°, and the water begins to 
boil, after which, no further change in temperature will 
occur until all the water has been changed into steam. 
[Another increase in the limit of molecular volume.] The 
heat which is required to boil the water is rendered latent. 
After the steam has been formed, it can be heated with 
increase of temperature and volume the same as any other 
gas. 

If the process is reversed and the gas is changed to a 
liquid or the liquid to a solid, the heat, which by the first 
process was rendered latent, now becomes sensible. This 
is the principle involved in heating buildings by steam. 

21. Heat Measurement. Sensible heat, or that which 
is shown by a rise of temperature, is measured by the 
thermometer, but latent heat cannot be measured in this 
way. For this purpose another unit must be employed, 
since it is evident that the thermometer cannot indicate 
the total amount of heat which is rendered latent. The 
total amount of heat employed in any process is measured 
by finding out how much water it is capable of heating 
from 0° to 1°. The amount of heat which is necessary to 
raise the temperature of one gram of water from 0° to 1° 
is called a calorie. This is the unit of heat. A certain 
amount of heat, say fifty calories, will raise one gram of 
water from 0° to 50° or will raise fifty grams of water 
from 0° to 1°. 

^ As there are different units of weight and different 
thermometric degrees it is evident that we could have 
different heat units, and indeed such are known. For all 
modern scientific measurements we use the calorie as 
defined above. If any other heat unit is employed, the 



18 LECTURE NOTES 

unit of weight and the thermometric degree is specified. 
An instrument for measuring the total amount of heat in 
a body, is called a calorimeter. 

22. Solutions. Solutions are homogeneous mixtures 
which cannot be separated by mechanical means. Solu- 
tions may be formed between substances in the same or 
in different states of aggregation. Two or more gases can 
always form a solution, and in all proportions, provided 
they do not unite chemically. Liquids very commonly 
form solutions with liquids, but not often in all propor- 
tions. The mixture of metals which we call alloys, may 
be considered as solutions of solids. 

The most important cases for our present consideration 
are the solution of solids in liquids and gases in liquids, 
water being the liquid used, unless otherwise stated. 

Nearly all solids are soluble in water to some extent, 
and, as a rule, the higher the temperature the more solu- 
ble the substance, although there are some exceptions to 
this rule. A solution is said to be saturated at a certain 
temperature if, when brought in contact with the sub- 
stance already in solution, no more will dissolve. A 
supersaturated solution is one which contains a greater 
amount of the solid substance than is required to form a. 
saturated solution. This condition can only be maintained 
when the solution is not in contact with the substance 
dissolved. When the liquid is removed from a solution 
by evaporation, the solid nearly always appears in the 
form of crystals. 

All gases dissolve in water to some extent and in some 
cases the amount is very great. The degree of solubility 
depends upon temperature and pressure, being increased 
by low temperature and by high pressure. The degree 



LECTURE NOTES 19 

of solubility for gases is usually reckoned in volumes, the 
total volume of the liquid being taken as the unit. If no 
chemical action takes place in the solution, gases can 
always be more or less completely expelled from the liquid 
by heating or by diminishing the pressure. 

23. Atoms. Atoms are the minute particles which by 
combination form the molecules and which cannot be 
divided into smaller parts by any chemical process. We 
may therefore define an atom as the ultimate division of 
the molecule, or, as the smallest particle of matter which 
can take part in a chemical change. 

Atoms may combine with other atoms of the same or 
of different kinds to form a molecule. If atoms of the 
same kind combine they form an elementary molecule, 
and a substance composed of elementary molecules is called 
an element. An element may be defined as a substance 
from which toe are unable to obtain any essentially different 
substance. 

There are at present about seventy-five different ele- 
ments known. 

When more than one kind of atoms combine they form 
a compound. From this fact it follows that a compound 
can always be decomposed into two or more different 

substances. 

24. Names and Symbols of the Elements. The names 
of the elements have been derived from various sources 
such as the names of mythological deities, planets, coun- 
tries and places, or from some property peculiar to the 
element or its compounds. Nine of the elements have been 
known since the very earliest times ; their discoverers 
and, in most cases, the origin of their names are unknown. 



20 LECTURE NOTES 

These elements are gold, silver, mercury, copper, iron, 
tin, lead, carbon and sulfur. 

Symbols are used for convenience in representing ele- 
ments. The symbol of an element is the initial letter of 
its name, to which there is added a second distinctive 
letter, when more than one element has the same initial. 
Ten of the elements, including those known to the ancients 
and some of those early discovered in modern times, have 
symbols taken from their Latin names. These are anti- 
mony, Sb, (stibium); copper, Cu, (cuprum); gold, Au, (au- 
rum) ; iron, Fe, (ferrum) ; lead, Pb, (plumbum) ; mercury, 
Hg, (hydrargyrum) ; potassium, K, (kalium) ; silver, Ag, 
(argentum); sodium, Na, (natrium); tin, Sn, (stannum); 
and tungsten, W, (wolframium). 

Each of these symbols, though sometimes used to rep- 
resent the name of the element, stands for a single atom 
of the element, and should always be so understood. 

The following is a complete list of the elements with 
their symbols and atomic weights : 



LECTURE NOTES 



21 



Name. 



Aluminum 
Antimony - 

Argon 

Arsenic -_ 
Barium — 
Bismuth __ 

Boron 

Bromin --_ 
Cadmium ._ 
Caesium — 
Calcium — 
Carbon _— 

Cerium 

Chlorin 

Chromium . 

Cobalt 

Columbium 

Copper 

Erbium — 
Fluorin — 
Gadolinium 
Gallium .-- 
Germanium 
Glucinum - 

Gold 

Helium 

Hydrogen 

Indium 

Iodin 

Iridium 

Iron 

Lanthanum 

Lead 

Lithium 
Magnesium 
Manganese 
Mercury 



Symbol. 



Atomic 
Weight. 



Al. 

Sb. 

A.(?) 

As. 

Ba. 

Bi. 

B. 

Br. 

Cd. 

Cs. 

Ca. 

C. 

Ce. 

CI. 

Cr. 

Co. 

Cb. 

Cu. 

Er. 

F. 

Gd. 

Ga. 

Ge. 

Gl. 

Au. 

He.(?) 

H. 

In. 

I. 

Ir. 

Fe. 

La. 

Pb. 

Li. 

Mg. 

Mn. 

Hg. 



27.1 
120.4 

40. (? 

75. 
137.4 
208.1 

11. 

80. 
112.4 
132.9 

40.1 

12. 
139. 

35.4 

52.1 

59. 

93.7 

63.6 
166. 

19. 
157. 

70. 

72.5 

9.1 

197.2 

4.(?) 

1.0075 

114. 

126.8 

193.1 

56. 

138.6 

206.9 

7. 

24.3 

55. 
200. 



Name. 



Molybdenum _. 
Neodymium 
Nickel 

Nitrogen 

Osmium 

Oxygen 

Palladium 

Phosphorus.-. 

Platinum 

Potassium- __. 
Praseodymium 

RHODIUxM 

Rubidium 

Ruthenium 

Samarium 

Scandium 

Selenium 

Silicon 

Silver 

Sodium 

Strontium 

Sulfur 

Tantalum 

Tellurium 

Terbium 

Thallium 

Thorium 

Thulium 

Tin 

Titanium 

Tungsten 

Uranium 

Vanadium 

Ytterbium 

Yttrium 

Zinc 

Zirconium 



Symbol. 



Mo. 

Nd. 

Ni. 

N. 

Os. 

O. 

Pd. 

P. 

Pt. 

K. 

Pr. 

Rh. 

Rb. 

Ru. 

Sm. 

Sc. 

Se. 

Si. 

Ag. 

Na. 

Sr. 

S. 

Ta. 

Te. 

Tb. 

Tl. 

Th. 

Tu. 

Sn. 

Ti. 

W. 

U. 

V. 

Yb. 

Yt. 

Zn. 

Zr. 



22 LECTURE NOTES 

25. Atomic Attraction. The atoms are held together 
in the molecule by a very strong attractive force which 
is known as chemical affinity, or chemism. 

This force differs entirely from gravitation or cohesion. 
It also varies greatly with the kind of atoms, but for the 
same kind of atoms under the same conditions, it is 
constant. 

Atoms of the same kind have some attraction for each 
other, as can be seen in the formation of elementary mol- 
ecules. This attraction differs very widely, and, in the 
case of a few metals, is very slight, as may be seen from 
the fact that their molecule consists of but a single atom. 

Atoms which have a similar chemical nature have little 
or no attraction for each other, but when they are of a 
widely different chemical nature the attraction between 
them is one of the strongest forces known. 

Chemical affinity can only manifest itself through mi- 
nute distances, so that any force which tends to increase 
the distances between the atoms weakens the attraction. 
As heat does this, chemical changes take place more 
readily at high than at low temperatures. 

26. Nascent Atoms. Atoms of the same kind have 
some attraction for each other since they combine to form 
elementary molecules. For this reason they do not ordi- 
narily exist alone. During the process of chemical change, 
atoms are constantly being liberated, and, for an instant 
of time, before they can combine with other atoms, they 
are free. While in this free condition they exhibit 
greatly intensified chemism, and are said to be in the 
nascent state. 

There are many chemical changes which take place 



LECTURE NOTES 23 

under the influence of these nascent atoms which cannot 
be accomplished by any other means. 

27. The Laws of Chemical Combination. When two 
or more elements combine to form a compound, they do 
so according to certain laws which govern all chemical 
combination. Experiment shows that the same kind and 
number of atoms is required to produce a given molecule, 
or, in other words, a given compound will always contain 
the same elements, combined in the same proportions by 
weight. This is called the law of definite proportions. 
This law was first clearly recognized and stated by Lavoi- 
sier, although not generally accepted until 1806, some 
years after his death. 

Long before an explanation of the fact was offered it 
had been noticed, that the same elements could sometimes 
be combined to form compounds which were entirely 
different. Among the elements which show this charac- 
teristic are oxygen and nitrogen, which form no less than 
five different compounds. About the year 1803, while 
experimenting with these compounds, Dalton discovered 
the fact, that when the same elements combine to form dif- 
ferent compounds, the different proportions by weight of a 
given element, will ahvays be simple mtdtiples of a common 
factor. This important discovery constitutes the law of 
multiple proportions. 

In the establishment of these two laws, many facts 
were brought to light. Among other things it was dis- 
covered that the relative proportion by weight of any 
element, which is required for a given compound, is cdways 
the same, or a midtiple of the same, id Inch is required for 
any other compound. This fact was discovered by Rich- 



24 LECTURE NOTES 

ter and is called the law of proportionality, or the law of 
reciprocal proportions. 

28. The Atomic Theory. The theory that matter is 
made up of minute particles (atoms) as opposed to the 
belief in the continuity of matter, that is, that matter 
completely fills the space it occupies, was taught by the 
Greek philosophers more than 2000 years ago. 

When Dalton discovered the law of multiple proportions, 
he proposed a more complete and scientific atomic theory 
to explain the laws of chemical combination. This theory 
has been universally accepted by scientists as the most 
important in all physical science. 

The most important principles which constitute the 
atomic theory are these : the ultimate constitution of all 
matter is atomic ; the atoms are unchangeable and 
chemically indivisible ; the atoms not only differ in kind, 
but the different kinds of atoms differ in weight ; the 
atomic weight is constant for a given kind of atoms ; and 
the changes which take place in matter, are due to a 
change in the kind, number, or relative position of the 
atoms which constitute the different kinds of matter. 

29. The Law of Gaseous Volumes. The laws of con- 
stant and multiple proportions, already stated, depend 
entirely upon weight relations. Soon after the atomic 
theory was proposed, G-ay-Lussac investigated the volu- 
metric relations of gases in chemical combinations. He 
found that when hydrogen and oxygen combined, it re- 
quired 2 volumes of hydrogen fori volume of oxygen, and 
that the result was 2 volumes of steam. (Water in the gas- 
eous form.) In a similar way he found that 1 volume of 
chlorin required 1 volume of hydrogen, forming thus 2 
volumes of hydrochloric acid ; that 1 volume of nitrogen 



LECTURE NOTES 25 

required 3 volumes of hydrogen, and that they produced 
2 volumes of ammonia. 

From these and many similar facts he derived the law 
of gaseous volumes, called also the law of Gay-Lussac, 
which is that when gases, either elementary or compound, 
unite chemically, the volumes of the gases required, bear a 
simple ratio to each other, and to the resulting compound 
in the gaseous state. 

It was the study of the facts embodied in this law that 
led Avogadro to his important hypothesis regarding mol- 
ecular volume. 

30. Atomic Weight. Previous to the time of Dalton's 
atomic theory, all atoms were supposed to have the same 
atomic weight ; but since that time it has come to be the 
universal belief, that the different kinds of atoms differ 
in weight. The absolute weights of the atoms can only 
be approximated, but their relative weights can be deter- 
mined with great accuracy. 

The natural unit of atomic weight is the atom of 
hydrogen, since that is the lightest atom known. Only 
a few of the elements combine with hydrogen directly 
and so it is necessary, in most cases, to employ some 
other element in determining atomic weight. The element 
most commonly employed is oxygen, with which nearly 
every known element will combine. For this reason, 
oxygen is often taken as the standard, its atomic weight 
being assumed as 16. Most of our tables of atomic weights 
are now made by this standard. 

The relative weights of hydrogen and oxygen have 
been very carefully determined, so that atomic weights, 
given by one standard, may be changed to those of the 



26 LECTURE NOTES 

other. If the atomic weight of hydrogen is 1, that of 
oxygen is 15.88, and if the atomic weight of oxygen is 
assumed to be 16, that of hydrogen is 1.0075.* 

31. Valence. This is the measure of quantitative 
atomic value. When two or more elements combine to 
form compounds, they do so according to the laws of 
chemism, or chemical affinity. The force which we call 
chemism is entirely qualitative. If we study different 
compounds we shall see that a given atom may require 
one, two, three, or even more atoms of some other kind 
to satisfy its chemism. This quantitative value in chem- 
ical combination we call valence. 

The unit of valence is the hydrogen atom. The hydro- 
gen atom is therefore univalent. Univalent atoms are 
such as require one atom of hydrogen, or of some other 
univalent element, to satisfy their chemism. Those atoms 
which require two univalent atoms are called bivalent ; 
those which require three or four univalent atoms, or 
their equivalent, are called trivalent or quadrivalent re- 
spectively, etc. 

A univalent atom is called a monad, atoms having a 
higher valence being called dyads, triads, tetrads, etc., 
respectively. 

The valence of an atom is not ordinarily indicated, but 
if, for any reason, it is desirable to indicate the valence 
of an atom, it may be done by means of a Roman numeral 
placed above the symbol ; or by short lines placed above 
or about the symbol, thus : 

I II in | IV 

CI, O, S = , N = , P , — C— or C , etc. 

I 
*For a table of atomic weights see list of the elements on 
page 21. 



LECTURE NOTES 27 

These marks, which only indicate the number of affinity 
units possessed by the atom, are quite often referred to, 
as bonds of affinity. 

When a compound is formed, all the affinity units, or 
bonds, of the combining atoms, must be satisfied. This 
can be done by any combination of atoms, possessing the 
equivalent number of affinity-bonds, subject, of course, to 
the laws governing atomic attraction. 

32. Variable Valence. It was at first believed that 
valence was a fixed quantity, something inherent in the 
atom ; but experiment shows that in certain combinations 
a given atom will vary in valence. 

The valence of an atom, therefore, seems to be depend- 
ent upon the nature of the atoms with which it combines ; 
for, when they combine with hydrogen the valence of 
atoms is found to be constant, but with oxygen, chlorin, 
and elements of a like nature, atoms exhibit their greatest 
variation in valence. 

Valence is sometimes modified by temperature since it 
is often seen that at a high temperature, an atom exhibits 
its lowest valence. 

In many compounds atoms are combined in a chain-like, 
or a ring-like union, in which some of the affinity bonds 
are satisfied by some of those belonging to atoms of the 
same kind. This is the explanation of the numerous 
compounds of carbon and hydrogen. 

The explanation, that the apparently lower valence is 
due to the formation of an unsaturated molecule, and that 
the higher valence is due to a molecular combination, is 
perhaps true for some combinations, but is not true for all. 



28 LECTURE NOTES 

The belief is therefore universal that some and perhaps 
most of the atoms may vary in valence. 

This variation in valence is undoubtedly governed by 
some general laws, but at present we do not know what 
these laws are. 

33. Classification of the Elements. The elements may 
be classified in several ways. The oldest and most common 
classification is the one which divides them into two 
general classes, called metals and non-metals. This clas- 
sification depends upon the nature of the physical and 
chemical properties which the elements possess. The 
principal physical properties regarded as metallic, are, 
hardness, high specific gravity, luster and conductivity 
of heat and electricity. A few of the metals lack some of 
these properties and some of the non-metals possess them. 

The principal distinguishing chemical property is, that 
the non-metals combine with hydrogen and oxygen to 
form a class of compounds which are called acids, while 
the metals combine in the same way to form an entirely 
different class of compounds which are called bases. This 
latter distinction would be perfect were it not for the fact 
that a few of the elements, which occupy an intermediate 
position, are so influenced by the nature of the elements 
with which they combine, that they act either as metals 
or as non-metals, as the case may be. There is, therefore, 
no very well defined line of separation between the two 
classes. 

The one gradually merges into the other, so that the 
division is quite arbitrary. In the list of the elements 
on page 21, the names of the non-metals are printed in 
small capitals. 



LECTURE NOTES 29 

For convenience of study and reference, the elements 
are often arranged in groups. These groups consist of 
those elements which have many common properties, or 
those which are found together in separating them by 
qualitative analysis. 

When the electric current is made to pass through a 
solution containing certain chemical compounds, decom- 
position takes place. Some of the atoms thus liberated, 
will appear at the positive pole and some at the negative. 
Those which appear at the negative pole are called electro- 
positive or positive atoms, and those at the positive pole, 
electro -negative or negative atoms. The terms positive 
and negative, as here used, are entirely relative ; for, 
while the metals, as a class, are positive, and the non- 
metals negative, we may say of the intermediate elements 
that they are positive in respect to the non-metals and 
negative in respect to the metals. The greater differences 
the atoms of a compound exhibit in their electrical affini- 
ties, the more powerful are their attractions for each 
other, and the more stable is the compound. 

34. The Periodic Classification. The most complete and 
logical classification is the one commonly known as the 
periodic system. If the elements, after hydrogen, are 
arranged in a table according to their increasing atomic 
weights, they will fall into periods, as they are called, and 
the corresponding elements in each period will form groups, 
in which the elements are seen to be closely related to 
each other. 

The belief that there was some relation between the 
properties of the elements and their atomic weights has 
been held for a long time, especially in the case of ele- 
ments which closely resemble each other in their chemical 



30 LECTURE NOTES 

affinities. This was enlarged and systematized in 1869 
by the Eussian chemist, Mendelejeff, who brought all the 
known elements into the system and formulated what has 
come to be known as the periodic law, which is that the 
properties of the elements, as well as the properties of their 
compounds, are periodic functions of their atomic weights. 

There are yet many questions unanswered by this 
classification, but the underlying principle is no longer 
questioned. It has been almost universally accepted as 
a fundamental law of chemistry. A more complete study 
of this classification, belongs in a more advanced course. 

35. Compound Radicals. There are certain groups of 
atoms which are found as an invariable constituent in 
a series of compounds. These groups can be replaced 
in their compounds by single elements, or by other 
groups. They do not often exist alone, although two 
such groups of the same kind may unite to form a com- 
pound, just as two elementary atoms may combine to form 
an elementary molecule. They therefore act just like a 
single atom. Such a group of atoms is called a compound 
radical. A compound radical is, therefore, a group of 
atoms which acts like a single atom. 

The compound radicals are classified according to their 
properties, and their relations to the elements, and in 
many ways are treated exactly as if they were elementary 
atoms. Some of them resemble the metals, while others 
are like the non-metals. 

Among the important compound radicals may be men- 
tioned ammonium, (NH 4 ) — , a monad group resembling 
the metals, sodium and potassium. Cyanogen, (CN) — , 
is a monad group which in its compounds, closely resem- 



LECTURE NOTES 31 

bles chlorin. A very important compound radical is the 
monad group called hydroxyl, (OH) — , which occurs in a 
multitude of compounds. It is sometimes combined with 
a metal, and sometimes with a non-metal, and imparts 
characteristic properties to each. 

36. Chemical Formulas. We have noted that each 
element can be represented by a symbol, which represents 
one atom of the element. In a similar way we can repre- 
sent each compound by a combination of symbols. Such 
a representation of a compound is called a molecular or 
chemical formula. 

There are two general ways of writing a chemical for- 
mula. The symbol for each atom may be repeated as 
many times as it occurs, or the number of each kind of 
atoms which occur in the compound, if more than one, 
may be indicated by a small numeral placed at the right 
and a little below each symbol. 

The latter method is the one commonly used to represent 
the less complicated compounds of inorganic chemistry. 
It shows only the ultimate composition of the molecule 
and is called an empirical formula. 

In representing the more complicated compounds, 
especially such as are often found in organic chemistry, a 
much clearer understanding of the nature of the compound 
may be obtained, if we indicate, in the formula, the mutual 
relations of the atoms and compound radicals. Such a 
formula is called a constitutional or rational formula. In 
this formula the valence of each atom is indicated by short 
lines which show how it is connected with the other atoms 
and how each valence bond is satisfied. 

As an illustration of these different formulas, we find 



32 LECTURE NOTES 

by analysis that the compound called sulfuric acid is com- 
posed of two parts of hydrogen, one part of sulfur and 
four parts of oxygen. 

We may express this fact by writing as its formula, 
H 2 S0 4 . By further investigation, we find that each 
hydrogen atom is very closely related to an atom of oxy- 
gen. The compound contains therefore two hydroxyl 
groups. For this and other reasons, we express these 
relations by the constitutional formula, 

H— O . Q # O 

H-0> S ^ - 
A large numeral placed before the whole or part of an 
empirical formula, denotes that the portion following the 
numeral is to be taken as many times as is indicated. 
Thus, 2 CuS0 4 means two molecules of CuS0 4 . A large 
numeral placed before or a small numeral placed after 
symbols or formulas placed in a parenthesis, affects all 
within the parenthesis. Thus Pb (N0 8 ) 2 , or Pb 2 (NO s ), 
means that the (N0 8 ) is to be used twice in the compound. 

37. Acids. We have already noted (33) that the non- 
metallic elements combine with hydrogen and oxygen to 
form compounds which are called acids. The non-metal 
combined in this way, imparts to the acid its own peculiar 
chemical qualities and so may be called the principal 
element in the acid. A compound radical, having the 
properties of a non-metal, may form an acid in the same 
way. All acids contain hydrogen, combined with the 
principal element or radical either directly or by means 
of oxygen. Among those which contain no oxygen, but 
have the hydrogen joined directly to the principal element, 
are the so-called halogen acids, while those containing 
oxygen are called oxy -acids. There are a few compara- 



LECTURE NOTES 33 

tively rare acids in which the oxygen is partly or entirely 
replaced by sulfur. These are called thio- or sulfo-acids. 

Acids are so Darned because they are generally sour to 
the taste. They also act upon vegetable colors, (such as 
litmus,) changing them from blue to red. This is the 
common test for an acid, and is called the acid reaction. 

The hydrogen of an acid may be partly or entirely 
replaced by a metal or positive radical, forming a com- 
pound which is called a salt. For many reasons, it is 
believed that the replaceable hydrogen in the oxy-acids is 
joined to the principal element by means of oxygen, that 
is, each atom forms a part of an hydroxyl group. Because 
of these facts this hydrogen is called basic hydrogen. 

Acids are further distinguished as mono- basic, di-basic, 
tri-basic, etc., according to the number of replaceable 
hydrogen atoms which they contain ; or, what is the same 
thing, the number of hydroxyl groups which they contain. 
This is called the basicity of an acid. 

Nearly all the non-metals combine with oxygen forming 
compounds which are called oxids, and which will combine 
with water, H 2 G, to form acids. Conversely, if we can 
remove the combined water from the acids we may obtain 
the oxid. Such oxids are sometimes called the anhydride 
of the corresponding acids. 

The group of atoms, which together with the replacea- 
ble hydrogen makes up the complete acid molecule, is 
called the acid radical. 

Thus sulfuric acid, H 2 S0 4 , is an oxy-acid containing two 
basic hydrogen atoms. It is therefore dibasic. If we 
remove water from it, there will be left its anhydrid, SO s . 



34 LECTURE NOTES 

The remainder after removing its basic hydrogen, is S0 4 , 
which is its acid radical. 

38. Bases. When a metal combines with hydrogen 
and oxygen, the compound formed is called a base. Bases 
also have peculiar properties, which are the opposite of 
those possessed by acids. The most characteristic prop- 
erty of a base, is its power to unite with an acid to form 
a neutral compound, called a salt. Only a few of the bases 
are soluble in water. The soluble bases are called alkalies, 
and they have a soft slippery feel and a peculiar soapy 
taste. They act upon vegetable colors, changing them 
from red to blue. This is the common test for a soluble 
base, and is called the basic or alkaline reaction. 

The compounds of some of the metals with oxygen alone, 
act in this way, and so they are sometimes called bases. 

For reasons similar to those given under acids, the 
hydrogen and oxygen in a base are supposed to exist in 
the form of an hydroxyl group. The hydrogen in a base 
is called acid hydrogen, and bases are distinguished as 
mon-acid, di-acid, tri-acid, etc. This is called the acidity 
of a base. In the same way that the hydrogen of an acid 
may be replaced by a basic radical, the hydrogen of a 
base may be replaced by an acid radical. In this action, 
the total acidity of the base must just equal the total 
basicity of the acid. A few bases consist of a group of 
atoms (one of them a metal) combined with hydroxyl. 

39. Salts. When an acid and a base combine with each 
other, they form a neutral body which is called a salt. A 
salt may be considered as derived from an acid, by replac- 
ing the hydrogen of the acid with the metal of the base ; 
or, as derived from a base, by replacing the hydrogen of 
the base, with an acid radical. As a base residue is simpler 



LECTURE NOTES 35 

than that of an acid, the former is the more common view, 
although the two are practically identical. Water is 
always formed in this process. A salt is, therefore, a 
compound, consisting of a metal, or positive radical, and 
a non-metal, or negative radical, combined usually by 
means of oxygen. 

The simplest conception of a salt is, that it is always 
derived from an acid by replacing the hydrogen with a 
metal or positive radical. All acids have a corresponding 
salt, and conversely, all salts are derived from some acid. 
Many salts are known which are perfectly stable, while 
the acids, from which they are derived, do not exist free. 

If the metal replaces all the hydrogen in the acid, which 
is capable of being replaced, the compound formed is 
called a normal salt. If apart of the replaceable hydrogen 
remains in the new compound, it is called an acid salt. 
Normal salts do not usually affect vegetable colors. They 
usually give, therefore, what is called a neutral reaction. 
Since acid salts retain some of their basic hydrogen, they 
usually give an acid reaction. 

We have already noted that in a few bases, a basic 
group of atoms is combined with hydroxyl. In such a 
compound the relative amount of metal to hydroxyl is 
greater than in an ordinary base. When such a base 
combines with an acid to form a salt, the proportional 
amount of metal is greater than in a normal salt. The 
same result obtains if we consider the new compound as 
derived from a base, in which only a part of the acid 
hydrogen is replaced by the acid radical. Such a com- 
pound is called a basic salt. 

Basic salts are usually insoluble in water, and may often 



36 LECTURE NOTES 

be recognized by this characteristic. Certain of the acids 
and bases show this tendency in a marked degree. Such 
are sulfuric, nitric and carbonic acids, and the bases of 
lead, mercury, copper, bismuth, antimony and zinc. 

In a few cases the normal salts give either an acid, or 
an alkaline, reaction. This depends upon the nature of 
the elements entering into combination. If a metal with 
very strongly metallic properties, combines with an acid 
which shows very weak chemical properties, the salt 
formed will give an alkaline reaction. Thus Na 2 CO s gives 
an alkaline reaction. On the other hand, a metal with 
very weak basic properties combined with an acid with 
very strong chemical properties, forms a salt which may 
show an acid reaction. For a similar reason an acid salt 
may show a neutral or even an alkaline reaction. 

40. Modes of Chemical Action. In order to effect a 
chemical change, it is necessary that the molecules come 
very closely together. This can only be secured when the 
molecules are able to move freely among themselves. For 
this reason the liquid or the gaseous state, is more favor- 
able to chemical change. The gaseous state is, theoreti- 
cally, most favorable, since in this condition the molecules 
move with the greatest freedom ; but because of the 
difficulty in obtaining many substances in this condition, 
and, because of the difficulty in handling gases, it is not 
generally employed, and so the liquid condition is, all 
things considered, best. Solids can be brought into the 
liquid condition either by solution or by fusion, and gases 
by solution. 

The degree of chemical activity differs greatly with the 
different elements. If an element, possessing powerful 
affinities (chemism), comes in contact with a compound, 



LECTURE NOTES 37 

it frequently happens that it displaces some element, or 
radical, in the compound, thus forming a new compound. 

There are several ways in which chemical action may 
take place. Two or more elementary, or simple molecules, 
may unite to form a more complex one. Such an action 
is called synthesis. 

A compound may be decomposed by some force so as to 
form two or more simpler compounds, or the compound 
may even be resolved into its elements. This action is 
called analysis. 

When two compounds come together, it often happens 
that the conditions are such as to bring about an exchange 
of their constituent parts. Such an action is called 
metathesis. This is by far the most common mode of 
chemical action. There are many causes which bring 
about the final result in metathesis. 

41. Chemical Equations. The changes which occur 
when molecules are brought together may be expressed 
by what is called a chemical equation, in which the formu- 
las of the molecules which are brought together, are placed 
before a sign of equality and called the factors, while the 
formulas of the resulting molecules, are placed after the 
sign of equality and called the products. 

In a chemical equation, every atom of the factors must 
be accounted for in the products ; and, since the atomic 
weights are constant, it follows that the sum of the mo- 
lecular weights of the products must equal the sum of the 
molecular weights of the factors. The following will serve 
as an illustration : 

AgNO s + NaCl = AgCl -f NaNO s 
This means that if a molecule of AgNO s and a molecule of 



38 LECTURE NOTES 

NaCl come together, there will result a molecule of AgCl 
and a molecule of NaN0 3 . Since the molecular weight of 
a substance is equal to the sum of the atomic weights of 
the atoms forming the molecule, if we substitute, in the 
above equation, the molecular weights of the compounds 
for their formulas, we have the following equation : 
169.9 + 58.4 = 143.3 + 85, or, 228.3 = 228.3. 
This shows that the sum of the molecular weights of 
the factors and of the products is the same. 

Since a chemical equation represents what actually takes 
place in a chemical change, we must know what this is 
before we can express it in an equation. We may deter- 
mine the products by a chemical analysis, and then, from 
our knowledge of the elements and of their relations to 
each other we can express what must have taken place. 
An example will make this clear. If we put a piece of 
copper into sulfuric acid (H 2 S0 4 ) and heat them, we get a 
gas, called sulfur dioxid (S0 2 ), and from the liquid we 
can get a compound called copper sulfate (CuS0 4 ). From 
the relations of the two compounds, CuS0 4 and H 2 S0 4 , 
we may safely conclude that the copper has replaced the 
hydrogen in the sulfuric acid. Expressing this by a 
formula we have: 

Cu + H 2 S0 4 = CuS0 4 + H 2 . 

This would indicate that hydrogen must be set free in 
this action ; but, as we get sulfur dioxid and not hydro- 
gen, the hydrogen must have had some action on the 
sulfuric acid, from which the sulfur dioxid must come. 
If we take away from H 2 S0 4 the atoms necessary to form 
sulfur dioxid, viz : S0 2 , we have left H 2 and 2 , which, 
with the H 2 set free by the copper, are in just the right 
proportions to form water, H 2 0. Since nothing else can 



LECTURE NOTES 39 

be found in the solution, we may express the result as 
follows : 

H 2 + H 2 S0 4 = 2H 2 + SO,. 

We may now combine the two equations, to express the 
complete result, as follows : 

Cu + 2 H 2 SO t = CuS0 4 + 2 H 2 -f S0 2 . 

42. Chemical Nomenclature. Until about the close of 
the eighteenth century, the number of elements and 
compounds was small and only arbitrary names were given 
to them. Compounds took their names from their discov- 
erer ; or from the locality in which they were discovered ; 
or from some physical property. Compounds which had 
some real or fancied resemblance were often grouped 
together. The first attempt at a systematic nomencla- 
ture, was made by Lavoisier, in 1787, and his system, 
with only slight modifications, is the one in use at the 
present time. A few of the old common names were 
preserved, and are still used commercially. 

No attempt was made to change the old names of the 
elements, but the names of all newly discovered metals 
end in turn, while the names of the non-metals end in on. 

Those substances which contain only two elements, or 
compound radicals, are called binary compounds. 

In binary compounds, the name of the metal, or positive 
radical, is written first. The name of the non-metal, or 
negative radical, follows, with the last one or two sylla- 
bles of the name changed to id. Should more than one 
compound exist between the same elements, the name of 
the metal receives a new termination. The name of the 
metal, in the compound which has the relatively greater 
amount of the negative, or non-metallic element, ends in 



40 LECTURE NOTES 

ic, the other ends in ous. Thus FeCl 3 is ferric chlorid and 
FeCl 2 is ferrous chlorid. 

The halogen acids are named from the elements which 
they contain, the names being a combination of those of 
both elements. The ending is ic. Thus HC1 is hydro- 
chloric acid. 

The names of the oxy-acids are derived from their 
principal elements, and end in ic or ous. These 
terminations have the same meaning as in binary com- 
pounds, the relative amount of oxygen determining the 
ending. If there are more than two acids which have the 
same principal element, the most common or important 
acid is the one ending in ic. Thus there are four oxy-acids 
of chlorin, viz : HCIO, H01O 2 , HC10 3 and HCIO^. The 
most important one is H01O 3 which is called, chloric acid, 
HC10 2 being therefore chlorous acid. The name of the 
first one, which contains less oxygen than chlorous acid, 
is obtained by prefixing hypo, (under). HCIO is therefore 
hypoch\orous acid. The name of the fourth one, which 
contains more oxygen than chloric acid, is obtained by 
prefixing hyper, or per, (over). HCIO^ is therefore hyper- 
chloric or peroh\oric acid. 

The bases are named as binary compounds, being 
regarded as hydroxids of the metals. Thus KOH is 
potassium hydroxid. 

A salt is always regarded as having been derived from 
an acid, and its name is therefore taken from the name 
of the acid from which the salt is derived. If the name of 
the acid ends in ic, the name of the salt will end in ate ; 
and if the name of the acid end in ous, the name of the 
salt will end in ite. If the name of the acid has a prefix, 



LECTURE NOTES 41 

the name of the salt will have the same prefix. The name 
of the metal contained in the salt remains unchanged, 
except in the case where it forms two salts with the same 
acid, when the name is changed by terminations the same 
as in binary compounds ; thus, a salt of sulfuric acid and 
iron, would in general be called iron sulfate ; but, since 
there are two compounds of iron and sulfuric acid, we call 
FeS0 4 , ferrous sulfate, and Fe 2 (S0 4 ) s , ferric sulfate. 

An exception is made in the naming of the salts derived 
from the halogen acids. Inasmuch as these are binary 
compounds, they are named in accordance with the rule 
for naming binary compounds. Thus NaCl, which is 
derived from hydrochloric acid, HC1, might have been 
called sodium hydrochlorate ; but as it is a binary com- 
pound it is called sodium chlorid. 

In order to know the acid from which a given salt was 
derived, replace the symbol of the metal in the salt with 
an equivalent number of hydrogen atoms. 



PART TWO 



The Non-Metals 



Hydrogen. Symbol, H. Atomic weight, i. Specific 
gravity, 0.0694. 

43. History and Occurrence. Hydrogen was first ob- 
served in the early part of the seventeenth century, 
although not at that time recognized as an element. It 
was first accurately described, and its elementary nature 
established, in 1766, by Cavendish, who called it inflamma- 
ble air. A few years later, Lavoisier showed its relations 
to oxygen in water, and gave it the name hydrogen. 

Hydrogen occurs free in the gases issuing from certain 
volcanoes, and, mixed with other gases, in natural gas. 
It has also been found as occluded gas in meteoric iron. 
By means of the spectroscope, it is found to exist free in 
the atmosphere of the sun and many of the fixed stars. 

Hydrogen occurs combined in water, in many of the 
rocks, and in all animal and vegetable compounds. 

44. Preparation. Hydrogen may be prepared in many 
ways, the most common method being, to replace it in some 
acid by means of some metal. Ordinarily we use zinc and 
sulfuric acid and the action is shown by the following 
equation : 

Zn + H,S0 4 = ZnS0 4 + H 2 . 



LECTURE NOTES 43 

Water may be decomposed into its element by means of 
the electric current, hydrogen appearing at the negative 
pole, thus : 

2 H 2 = 2 H 2 + 2 
If certain metals, such as sodium or potassium, are 
brought in contact with water at the ordinary temperature, 
a part of the hydrogen will be replaced by the metal and 
come off free, thus : 

Na 2 -f 2 H 2 = 2 NaOH + H 2 . 
Other elements, such as zinc, iron, or carbon, will de- 
compose water at a high temperature giving hydrogen, 
thus : 

Zn + H 2 = ZnO -J- H 2 . 

45. Properties. Hydrogen is a gas. When pure, it is 
colorless, odorless, and tasteless. It is the lightest kind 
of matter known, and is 14.39 times as light as the air. 
A liter of hydrogen at 0° and 760 m. m. pressure weighs 
0.0898 grams. This is called a crith. When subjected to 
very great cold and pressure, hydrogen may be condensed 
to a colorless liquid. It is only very slightly soluble in 
water. 

Hydrogen is an inflammable gas and, when pure, burns 
with an almost colorless flame, which gives very little 
light, but a great amount of heat. When hydrogen burns, 
it does so by combining with the oxygen in the air, and 
forms water, thus : 

2H ! + 2 = 2H 2 0. 

Hydrogen has great affinity for the elements oxygen 
and chlorin, and in the decomposition of compounds con- 
taining these elements when hydrogen is present, water 
or hydrochloric acid is always produced. It forms com- 
pounds with all the non-metallic elements. 



44 LECTURE NOTES 

Hydrogen does not support combustion, and a lighted 
taper when plunged into the gas is at once extinguished. 

Pure hydrogen is not poisonous, and may be breathed 
for a short time without injury ; long continued breathing 
would cause death by suffocation. It produces a curious 
effect upon the voice, weakening it and elevating its 
pitch. 

Hydrogen will diffuse itself through certain red hot 
metals, such as iron, platinum, and palladium. These 
metals will retain a certain amount of the gas when 
cooled, the metal palladium retaining 935 volumes. This 
property has been termed occlusion by Graham, and, for 
various reasons, it is supposed that the hydrogen exists 
in them in the solid state, forming thus a kind of alloy 
with, the palladium. The name hydrogenium has been given 
to this form of hydrogen. 

Oxygen. Symbol, 0. Atomic weight, 16. 

46. History and Occurrence. Oxygen was discovered 
in 1774 by Priestley, who obtained it by heating mercuric 
oxid. A little later it was discovered independently by 
Scheele. Priestley called it dephlogisticated air ; the name 
oxygen was given to it later by Lavoisier. 

Oxygen is the most abundant element in nature. It 
occurs free in the atmosphere, of which it forms about 23 
per cent, by weight. Combined with hydrogen, it forms 
water, which is 88.81 per cent, oxygen. Combined with 
other elements, it is found in nearly all animal and vege- 
table matter, and constitutes about 48 per cent, of the 
earth's crust. The whole earth, including the surround- 
ing atmosphere, is about 50 per cent, oxygen. 

47. Preparation. Oxygen may be prepared from certain 



LECTURE NOTES 45 

of its compounds with the metals, called oxids, and also 
from its more complex compounds. 

The best and most common source of oxygen is potas- 
sium chlorate, KC10 3 . When this salt is heated, it loses 
all of its oxygen, leaving potassium chlorid, KC1, thus: 

2 KC10 3 = 2 KC1 + 3 2 . 
This action does not take place until the temperature 
is quite high, considerably above the melting point of the 
salt (359°), and unless care is taken, the decomposition 
may go on so rapidly as to cause an explosion. If about 
10 per cent, of manganese dioxid, Mn0 2 , is mixed with 
the potassium chlorate before heating, the oxygen will be 
given off before the salt fuses and there is very little 
danger of explosion. After the reaction, the manganese 
dioxid is found unchanged in composition, mixed with the 
potassium chlorid. 

The peculiar force exerted by the manganese dioxid in 
this action is called catalysis, or contact action. No satis- 
factory explanation of catalytic action has been given. 

When mercuric oxid, HgO, is heated, it is completely 
decomposed into mercury and oxygen, thus : 

2 HgO = 2 Hg + O,. 

The oxids of silver, gold and similar metals, act in a 
similar way. 

Certain of the oxids, which are called peroxids, give off 
a portion of their oxygen when heated. Manganese 
dioxid, Mn0 2 , yields one-third of its oxygen leaving man- 
gano -manganic oxid, Mn 3 4 , thus : 

3 Mn0 2 = Md 3 4 + 2 . 

Lead dioxid, PbO a , gives off one-half of its oxygen, 
leaving lead oxid, PbO. 



46 LECTURE NOTES 

48. Properties. Oxygen is a gas. It is colorless, 
odorless, and tasteless. It is slightly heavier than the 
air, its specific gravity being 1.104. It is only very 
slightly soluble in water. When it is subjected to very 
great cold and pressure, it becomes a bluish transparent 
liquid. 

Oxygen forms at least one compound with every known 
element, except fluorin. These compounds are called 
oxids, and many of them can be obtained by direct union 
of the elements. 

Oxygen does not burn, but it supports the combustion 
of inflammable substances with great vigor. It is generally 
recognized, and distinguished from other gases, by this 
property ; for, a splinter of wood, with a glowing spark 
upon it, immediately bursts into flame when thrust into 
this gas. [Another gas, nitrous oxid, also supports 
combustion, but may be distinguished from oxygen by 
other means.] All substances which burn in the air will 
burn with much greater rapidity and increased brilliancy 
in oxygen. 

Combustion, in the ordinary use of the term, is union 
with oxygen which is attended with light and heat. As 
the products of combustion are oxids, combustion is often 
called oxidation. 

Oxygen is necessary to all animal life. In the process 
of respiration the blood is purified by oxidation, and the 
waste products are thrown off by the lungs. The decay 
of animal and vegetable substances is caused by oxidation, 
or slow combustion. 

49. Oxidizing and Reducing Agents. An oxidizing 
agent is a substance which causes an element or compound 



LECTURE NOTES 47 

to combine with oxygen, or with other elements of a 
similar nature. When it is possible, oxidation changes 
an element from, a lower to a higher state of valence. The 
most common oxidizing agents are, oxygen, chlorin, 
potassium chlorate, nitric acid, and hydrogen dioxid. 

A reducing agent acts in an exactly opposite way, and 
so removes oxygen, or elements of a similar nature, from 
compounds. When it is possible, reduction changes an 
element from a higher to its lowest state of valence. It 
is the process of reduction which is employed in getting 
metals from their ores. The most common reducing agents 
are, hydrogen, (particularly in the nascent state,) carbon, 
sulphur dioxid, and hydrogen sulfid. 

A clear understanding of these two actions is very 
important to a student of chemistry. 

Ozone (Allotropic Oxygen.) Molecular formula, 3 . 
Density, 24. 

50. History and Occurrence. The peculiar odor which 
is found in the vicinity of an electric machine, when in 
use, and which follows an electric discharge through 
oxygen, was for a long time a puzzle to chemists. In 1840, 
this phenomenon was thoroughly investigated by Schon- 
bein, who showed that it was caused by a peculiar form 
of oxygen, which he called ozone. 

Ozone is an allotropic form of oxygen. When a sub- 
stance exists under two or more physical modifications, 
these are called allotropic forms. The allotropy in the 
case of ozone, is due to the fact, that its molecule contains 
three atoms. 

Ozone is generally believed to exist in very small quan- 
tities in the atmosphere, where it is produced by the 



48 LECTURE NOTES 

electric discharges during a thunder storm, also, proba- 
bly, by the growth of plants, and the lashing of waves. Its 
actual existence in the atmosphere is not easy to prove, 
since certain other substances, known to exist therein, 
have somewhat similiar properties. As it is a very ener- 
getic oxidizing agent, it is found most where the air is 
purest, as in the country, on the sea, and at high alti- 
tudes. 

51. Preparation. Ozone is best prepared by the action 
of a silent electric discharge on oxygen. An instrument 
which is used for preparing it in this way is called an 
ozonometer. It is also formed by the discharges of an 
electric machine through air or oxygen. 

It is evolved at the positive pole, together with oxygen, 
in the electrolysis of water. 

It is formed by the slow oxidation of phosphorus, and 
many other substances, and by the decomposition of 
certain highly oxidized compounds by means of sulfuric 
acid. 

All freshly prepared oxygen, from whatever source, 
contains more or less ozone. 

52. Properties. Ozone is a gas, having a density just 
one and a half times that of oxygen. It has a bluish 
color and a peculiar odor, a little like very dilute chlorin. 
Chemically, it is oxygen, but oxygen with greatly inten- 
sified properties. It oxidizes mercury, and silver in the 
cold. It liberates iodin from its compounds, and it is 
this action that is usually employed in detecting its pres- 
ence. It also oxidizes all organic substances. 

Only about five or six per cent of oxygen can be chang- 
ed into ozone, because the ozone xnolecules decompose 



LECTURE NOTES 49 

when brought into too close proximity with each other. 
At the ordinary temperature ozone decomposes slowly, 
and at about 300° it decomposes instantly, forming ordi- 
nary oxygen. 

COMPOUNDS OF OXYGEN WITH HYDROGEN 

There are two compounds of these elements, 
Hydrogen monoxid, or water, H 2 0, and 
Hydrogen dioxid, or peroxid, H 2 2 . 

Water. Molecular Formula, H ? 0. Molecular Weight, 18. 
Density, (of steam), 9. Specific Gravity, 1. 

53. History and Occurrence. Water was supposed, by 
the ancients and the alchemists, to be an element, and 
it was not until 1776, that Lavoisier proved it to be a 
compound. Five years later, Cavendish proved it to be 
composed of hydrogen and oxygen only, and in 1805, G-ay- 
Lussac proved the volumetric relations of its constituents. 

Water is the most abundant compound in nature. It 
is found free in the atmosphere, the amount being depend- 
ent upon the temperature, and forming a relatively small, 
but important part of that body. It is found in large 
quantities, both upon the surface and within the crust of 
the earth. As water of crystallization, it forms a neces- 
sary part of many mineral and chemical compounds. It 
is found in all animal and vegetable compounds, being- 
necessary to all animal and plant life and growth. 

54. Preparation. Water may be prepared by the direct 
union of its elements, that is, synthetically, by burning 
hydrogen in the air or in oxygen. 

When the two gases are mixed in the proportion of two 
volumes of hydrogen to one of oxygen, they can be made 



50 LECTURE NOTES 

to unite by a flame, or an electric spark, or by heating to 
about 600°, and the union is attended with violent explo- 
sion. If hydrogen is burned in a jet with oxygen, the 
heat is the most intense which can be obtained by com- 
bustion. 

Water can be prepared by the oxidation of compounds 
containing hydrogen, and hence it is always found as a 
product of the decay, or burning, of all animal or vege- 
table substances, and is thrown off by the lungs in the 
process of respiration. 

Many of the oxids of the metals when heated in hydro- 
gen, reduce to the metal and form water, thus : 
CuO + H 2 = Cu -f H 2 0. 

55. Properties. Water is a liquid, at the ordinary 
temperature, and, when pure, has neither odor nor taste. 
It is apparently colorless, although in large quantities, it 
has a distinctly greenish blue color. 

It has no action upon vegetable colors, and is a poor 
conductor of heat, and a worse conductor of electricity. 
Water is the most valuable solvent known. It not only 
dissolves many solids, but many liquids mix with it in all 
proportions. It dissolves all gases to some extent, and 
many in large proportions. On this account water s is 
never found quite pure in nature. The natural waters 
are usually pure enough for most purposes, and if pure 
water is required, it can be obtained by the process of 
distillation. 

Pure water is taken as the standard for the specific 
gravity of solids and liquids, and is used to obtain the 
unit of absolute weight in the metric system. 

Water forms a curious and important exception to the 



LECTURE NOTES 51 

general law that matter expands by heat ; for, if water is 
heated from 0° to 4°, it is found not to expand, but to 
contract. Above 4° it begins to expand and continues to 
do so, although the amount of the expansion varies with 
the temperature and is irregular. This peculiarity in the 
expansion and contraction of water is expressed by saying, 
that the point of maximum density of water is Jf°. 

Although the amount of contraction from 0° to 4° is 
very small, it is one of the most important facts in nature. 
It is for this reason that water freezes on the surface, 
and not, as is the case with most liquids, on the bottom. 
If this were not so, our climate would soon become frigid ; 
for our lakes and rivers would freeze solid in winter, and 
the summer heat would not suffice to melt the ice. 

Water usually becomes a solid, that is, freezes, at 0°. 
At the moment of becoming solid it increases about ^ I 
in bulk, the specific gravity of ice being 0.916. This is 
another important fact in nature, and is the principal 
cause of the disintegration of the rocks, and the forma- 
tion of soils. 

If water is carefully cooled, the temperature may often 
be reduced to some degrees below zero before the water 
is frozen, so that the freezing point of water is not neces- 
sarily 0°. It is for this reason that the melting point of 
ice, which is always 0°, is taken as one of the fixed points 
on the thermometric scale. 

When water is heated to 100°, under a barometric 
pressure of 760 millimeters, it is said to boil, and in 
boiling it forms a gas which we call steam. Above the 
temperature at which water boils, steam acts like a true 
gas, and obeys all the laws of gases. 



52 LECTURE NOTES 

56. Latent Heat of Water and Steam. We have already 
noticed, (20), that when ice is melted or water boiled, 
much heat is rendered latent. The amount of heat thus 
absorbed, or rendered lament, can be determined as follows : 
If we take 1 kilogram of water at the temperature of 0°, 
and 1 kilogram of water at 79°, and mix them, we shall find 
that the temperature of the mixture will be 39.5°; but if 
we take 1 kilogram of ice at 0°, and mix with it a kilogram 
of water at 79°, we shall find that the whole of the ice is 
melted, but the temperature of the water is exactly 0°. 
It has, therefore, taken the 79° of heat in the hot water 
to melt the ice, or, it requires 79 thermal units, (21), to 
melt ice. We say, therefore, that the latent heat of water 
is 79 thermal units, or calories. 

Under the normal barometric pressure water boils at 
100°. If we take 1 kilogram of water at 0°, and lead 
into it steam from boiling water, having the temperature 
of 100°, until the cold water has been made to boil, we 
shall find that the total weight of water is now 1.187 
kilograms, or 1187 grams. This means that the heat 
given off by the condensation of 187 grams of steam, has 
raised 1000 grams of water from 0° to 100°. To find the 
number of heat units, we must find what 1 kilogram, or 
1000 grams, of steam will do, which we get from the pro- 
portion, 187 : 1000 :: 1 : 5.36. If, therefore, 187 grams 
of steam will heat 1 kilogram of water from 0° to 100°, 
1000 grams, (or 1 kilogram) will heat 5.36 kilograms of 
water from 0° to 100°, or, what is the same thing, 536 
kilograms from 0° to 1°. The latent heat of steam is, 
therefore, 536 thermal units, or calories. 

The total amount of heat which is rendered latent in 
the formation of steam is given off when the steam con- 



LECTURE NOTES 53 

denses to water ; and since there are 536 calories to be 
liberated, it is easy to to see why steam is so effective in 
heating buildings. 

57. The Tension of Aqueous Vapor. It* we allow a 

glass of water to stand in a room for a number of days, 
the water will disappear, or evaporate, as we say. Even 
ice, if exposed to the atmosphere at a temperature below 
the freezing point, will evaporate. This power of water 
to pass into the form of vapor at all temperatures, is 
called the elastic force, or tension, of aqueous vapor. The 
vapor, which is thus given off, exerts a certain amount of 
pressure, which increases with the temperature. We can 
measure the amount of this pressure by placing a few 
drops of water above the mercury in a barometer, since 
the tension of the water vapor will depress the column of 
mercury. If the water is heated, the column of mercury 
will gradually sink until the temperature of the water 
reaches 100°, when it will be found to stand at the same 
level as that in the cup outside, showing that the tension 
is just equal to the atmospheric pressure. 

It follows, therefore, that water boils -when the\tension of 
its vapor is equal to the atmospheric, pressure. From this 
it follows that the temperature of the boiling point depends 
upon the atmospheric pressure. At an elevation high 
above the sea level, as on a mountain, the pressure is less 
and the boiling point is lowered, while in a steam boiler, 
where the pressure is increased, the boiling point is raised. 
The tension is measured in millimeters of mercury. [A 
table showing the tension for different temperatures will 
be found in most large works on chemistry.] 

In reckoning out the change in the volume of a gas, due 
to a change in temperature and pressure, if the gas is 



54 LECTURE NOTES 

measured over water, the tension of aqueous vapor must 
be taken into account. 

58. Water of Crystallization. When the water is 
evaporated from the solution of a salt, the salt nearly 
always takes a crystalline form. In a great number of 
cases it is found that a certain definite quantity of water 
is necessary to the formation of the crystals ; for when 
the water is removed, as it may be by heat, the crystals no 
longer retain their shape, but fall to powder. The water 
thus combined is called water of crystallization. 

The amount of water found in the crystals is reckoned 
in molecules, comparison being made with a single mole- 
cule of the salt. As many as ten or twelve molecules, or 
even more, are quite often found in a crystal. The 
amount which a given compound requires for crystalliza- 
tion often varies with the temperature of the solution, 
but for a given temperature it is invariable. 

The temperature at which crystals part with their water 
of crystallization is quite variable. Some lose a part of 
their water on exposure to the air, the water slowly evap- 
orating, and the salt falling to powder. This giving up 
of water at the ordinary temperature is called efflorescence. 
Some salts lose a part of their water at 100°, and require 
a higher temperature, before the rest will pass off. Thus 
copper sulfate, CuS0 4 , crystallizes with five molecules of 
water. Four of these molecules of water pass off at 100°, 
the fifth requires a temperature of about 250° to drive it 
off. The change of color in some of the so-called sympa- 
thetic inks, is caused by the giving up of water when they 
are heated. 

There are other crystals which combine with water when 



LECTURE NOTES 55 

exposed to the air and become partly or entirely liquified 
by the absorption of moisture from the air. This taking 
on of water from the air is called deliquescence. This 
property of certain salts is utilized in the construction of 
desiccators and desiccating tubes, for drying different 
substances. 

59. Natural Waters. Since water dissolves so many 
substances, the natural waters are never entirely free 
from impurities. These impurities may consist in part of 
insoluble matter, held in suspension by the water, and 
easily removable by nitration. The soluble matter can 
only be removed by distillation, or treatment with chemi- 
cals. Rain water is the purest form of natural water, 
and this contains dust and soluble impurities, washed 
from the atmosphere. Spring and river water always 
contains substances dissolved from the soil through which 
the water flows. 

When water contains an insufficient quantity of dis- 
solved matter to affect its taste, it is termed fresh water ; 
whereas, when it has a distinct taste, or possesses medic- 
inal qualities, it is called mineral toater. Those salts which 
are most often found in water, are the carbonates, sulfates 
and chlorids of calcium, magnesium and sodium, and 
occasionally salts of potassium, lithium and iron. Other 
compounds are also found, some of which are gases. 
Among the latter are carbon dioxid and hydrogen sulfid. 

Water is called hard, or soft, according as it contains 
large or small quantities of calcium or magnesium salts. 
The hardness of water is shown by its action upon soap, 
and determined by the amount of the latter which is nec- 
essary to produce a permanent lather in the water. Some 
salts may be removed from water by boiling. Such salts 



56 LECTURE NOTES 

produce what is called temporary hardness, while those 
salts which cannot be removed by boiling produce perma- 
nent hardness. 

Spring and river water, especially that which flows 
through wooded districts, contains more or less vegetable 
matter in solution. Such water is sometimes used for 
drinking, and, while not to be recommended for such 
purposes, its use is attended with comparatively little 
harm. The water supply which comes from shallow wells, 
is not infrequently contaminated by substances which 
come from the decomposition of animal matter. This is 
also true of water taken from rivers into which sewage 
enters above the intake, unless care is taken to remove 
this impurity. Such water should at once be condemned 
for drinking purposes, since it is liable to contain the 
bacteria of infectious diseases, and there is no longer any 
doubt that a number of infectious diseases are contracted 
and spread by means of impure drinking water. 

The contamination which exists in water can generally 
be removed by mechanical or chemical means, and the 
water rendered fairly safe for use. Water can be freed 
from bacteria by long continued boiling, and by properly 
constructed filter beds. 

Hydrogen Dioxid. Molecular Formula, H 2 2 . Molecu- 
lar Weight, 34. 

60. Occurrence and Preparation. This substance was 
discovered in 1818 by Thenard. It occurs in minute 
traces in the atmosphere, and in rain water and snow. 

It is best prepared by decomposing pure barium dioxid 
with dilute sulfuric acid, thus : 

Ba0 2 + H 2 S0 4 = BaS0 4 + H 2 2 . 



LECTURE NOTES 57 

The barium dioxid is slowly added to the dilute acid 
(1 :5) until the mixture is only slightly acid, the insoluble 
barium sulfate allowed to settle and the liquid filtered off 
and concentrated in a vacuum over sulfuric acid. It can 
also be prepared by dissolving sodium dioxid, Na 2 2 , in 
water and adding a dilute acid. 

61. Properties and Uses. Hydrogen dioxid, (called 
also hydrogen peroxid), is a colorless, transparent, oily 
liquid, having a specific gravity of 1.45. It has a bitter 
taste, and when placed upon the skin produces a white 
blister. 

It is a very unstable compound, easily decomposing into 
oxygen and water. At 100° the concentrated solution 
decomposes so rapidly as sometimes to give rise to an 
explosion. A dilute solution of this compound in water 
is much more stable, especially if it contains a drop or 
two of sulfuric acid. It dissolves easily in ether, the 
etherial solution being much more stable than the aqueous. 

It is used for bleaching silk and wool and for bleaching 
the hair ; also for restoring the colors of blackened oil 
paintings, and for removing stains from old manuscripts 
and engravings. In analytical chemistry it is used as an 
oxidizing agent. 

62. The Halogen Group. The four elements, fluorin, 
chlorin, bromin and iodin, constitute a group known as 
the halogens, or salt producers. These elements differ 
considerably in their physical properties, fluorin and 
chlorin being gases, bromin a liquid, and iodin a solid. 
They are, however, very much alike in their chemical 
properties, and possess very powerful chemical affinities. 
They all combine with hydrogen to form the halogen acids, 



58 LECTURE NOTES 

and with most of the metals, directly, to form salts. They 
are all univalent in their compounds with hydrogen, and 
the metals, but in compounds with oxygen they often 
exhibit a higher valence. The chemical affinity of these 
elements varies inversely as their atomic weights, fluorin 
having the strongest and iodin the weakest affinity of the 
four. Chlorin is the most common, and the typical 
element of the group. 

The compound radical cyanogen, (CN), acts much like 
the elements of this group. It forms compounds in every 
way analogous, and so is often classed as a member of the 
group. Its salts are called cyanids. 

Chlorin. Symbol, CI. Atomic Weight, 35.4. 

63. History and Occurrence. Chlorin was first obtained 
in the free state in 1774 by Scheele. who prepared it by 
the action of hydrochloric acid upon an ore of manganese. 
He at first supposed it to be an oxid of hydrochloric acid, 
and so called it deyj/ilogtisticated marine acid air. In 
1810, Davy proved it to be an element, and gave it its 
name, chlorin. 

Chlorin never occurs free in nature, on account of its 
strong affinity for most of the metals. Its most important 
compound is sodium chlorid, or common salt, which is 
found in large quantities in sea- water, in salt lakes, and 
in mineral springs. The chlorids of magnesium, calcium, 
and potassium are always found associated with the sodium 
chlorid. 

Sodium chlorid also occurs, in the form called rock-salt, 
in large deposits which are found in various parts of the 
earth. The most important of these deposits are atStass- 



LECTURE NOTES 59 

furt, in G-ermany,in the Turk's islands, W. I., in Louisiana 
and in western New York. 

64. Preparation. Chlorin is most conveniently pre- 
pared by the action of hydrochloric acid upon manganese 
dioxid, thus : 

4 HC1 -f Mn0 2 = MnCl 2 -f 2 H 2 + Cl 2 

It can also be obtained from its most common compound, 
sodium chlorid, by heating this with sulfuric acid and 
manganese dioxid, thus : 

2 NaCl 4- 2 H SO A + MnO, = 

MnS0 4 -f Na 2 S0 4 -f 2 H 2 -f Cl 2 . 

Almost any highly oxidized compound, if heated with 
hydrochloric acid, will yield chlorin. 

There are several processes for the manufacture of 
chlorin on a large scale, the principle involved being in 
all cases similar to that in the equations just given. 

65. Properties. Chlorin is a gas, having a yellowish- 
green color, and a most disagreeable and suffocating odor, 
which, when present in very small quantities, resembles 
that of seaweed. 

If chlorin is breathed, even in small quantities, it acts 
as an irritant poison, producing coughing and inflamma- 
tion of the mucous membrane ; if inhaled in the pure state, 
it may even cause death. It is about two and one-half 
times as heavy as the air, and a liter of the gas at normal 
temperature and pressure weighs 3. 18 grams. Under a 
pressure of six atmospheres at 0°, it becomes a yellow 
liquid. It is quite soluble in water, one volume of which 
dissolves about 2.5 volumes of chlorin at the ordinary 
temperature. If a solution in water is exposed to sun- 



60 LECTURE NOTES 

light, the water will be decomposed, hydrochloric acid and 
oxygen being formed, thus : 

2H 2 0+2Cl 2 = 4HCl+0 2 . 

Chlorin combines directly with all the common elements 
except oxygen, nitrogen, and carbon, and it combines 
with these elements indirectly. Many of the metals com- 
bine with chlorin at the ordinary temperature, giving 
light and heat, and forming a chlorid, while others require 
heating before combination will take place. 

Chlorin has a very strong affinity for hydrogen, as seen 
in its decomposition of water, and so it is a strong oxidiz- 
ing agent. In the presence of water it destroys organic 
coloring matter and so is much used in the process called 
bleaching. 

The oxidizing action of chlorin makes it valuable for the 
destruction of the poisonous germs of disease, and for 
removing the bad odors arising from the decay of organic 
matter. It is on this account much used as a disinfectant. 

66. Bleaching. Bleaching is the process of removing 
the color from a substance, so as to make it white, or 
light colored. Chlorin will bleach only in the presence 
of water. In the operation hydrochloric acid is formed, 
and oxygen is liberated, hence the action is one of oxida- 
tion, the color being destroyed by the liberated oxygen. 

Bleaching is a very important industry in the arts, espec- 
ially in the manufacture of cotton and paper. The excess 
of chlorin, and the hydrochloric acid, produced in the 
operation, would not only damage the machinery, but 
would injure the strength of the material bleached, and so 
they have to be removed. The chlorin is removed by the 
addition of what is called by the paper makers an antichlore, 



LECTURE NOTES 61 

and sodium hyposulfite is generally used for this purpose. 
The acid is removed by the addition of some weak alkaline 
salt, such as sodium carbonate. Chlorin will not usually 
bleach mineral colors, or those produced by carbon, so 
that printer's ink is not affected by chlorin. 

COMPOUND OF CHLORIN WITH HYDROGEN 

Hydrogen Chlorid, or Hydrochloric Acid. Molecular 
Formula, HC1. Molecular Weight, 36.4. Density, 
18.2. 

67. History and Occurrence. Hydrochloric acid, known 
to the Arabian alchemists under the name " spirit of salt,''' 
was known to most of the alchemists, in solution. It was 
first prepared as a gas, by Priestley, who collected it over 
mercury and called it marine acid air. It was supposed 
to be a compound of oxygen until 1810, when its true 
composition was determined by Davy. 

It occurs in the gases from active volcanoes, and in the 
water of certain rivers which flow through volcanic dis- 
tricts. Its salts are called chlorids, and occur as given 
under chlorin. 

68. Preparation. The common way to prepare hydro- 
chloric acid, is by acting upon some chlorid with sulfuric 
acid. Sodium chlorid is usually employed, and the fol- 
lowing equation represents the action : 

2 NaCl + H 2 S0 4 = Na 2 S0 4 -f 2 HC1. 

If the sulfuric acid is in excess of the sodium chlorid, 
acid sodium sulfate is formed with the hydrochloric acid, 
thus : 

NaCl -f H 2 S0 4 = NaHS0 4 -f HC1. 

Hydrochloric acid is also formed by the direct union of 



62 LECTURE NOTES 

its elements, that is to say, by synthesis. Hydrogen and 
chlorin will combine by heat, by the electric spark, or by 
the action of sunlight at the ordinary temperature, the 
union taking place with explosive violence. 

Commercial hydrochloric acid is obtained in enormous 
quantities, as a by product in the manufacture of sodium 
carbonate, by the Leblanc process. 

69. Properties. Hydrochloric acid is a colorless gas. 
It has a pungent, suffocating odor, and a sour taste. 
Under strong pressure at a low temperature, it con- 
denses to a colorless liquid. It is a little heavier than 
air, its specific gravity being 1.27. It does not burn nor 
does it support combustion. It fumes strongly in ordi- 
nary air, owing to its union with atmospheric moisture. 

It is very soluble in water. One volume of water at 0° 
dissolves about 500 volumes of the gas, and at the ordi- 
nary temperature about 450 volumes. This solution is the 
form in which hydrochloric acid is ordinarily used in the 
laboratory, and in the arts. It is known, commercially, 
as muriatic acid. 

When this aqueous solution is saturated, it contains 
about 42 per cent, of hydrochloric acid, and has a specific 
gravity of 1.21. When such a solution is heated, the 
gas is given off until the strength has been reduced to 
about 20 per cent., and the specific gravity to 1.10. At 
this strength it boils without further change. 

COMPOUNDS OF CHLORIN WITH OXYGEN AND 
HYDROGEN 

70. The Oxids of Chlorin. Chlorin forms two oxids, 

Chlorin monoxid, C1 2 0, and 
Chlorin peroxid, C10 2 . 



LECTURE NOTES 63 

Both of these compounds are gases, at the ordinary 
temperature. They are very unstable, and a very slight 
elevation of temperature causes them to explode with 
considerable violence. The monoxid has no practical 
value. 

71. Chlorin Peroxid. This compound is formed when 
a chlorate is decomposed by sulfuric acid, at a little above 
the ordinary temperature. Potassium chlorate is usually 
employed for this purpose, and the reaction is as follows: 

3KC10 3 + 3H 2 S0 4 r=: 3KHS0 4 + HC10 4 -f H 2 + 201O 2 . 

The action is somewhat more complicated than appears 
in the equation, chloric acid, HC10 3 , being first formed. 
This acid is not very stable and so decomposes at a slight 
elevation of the temperature, forming the chlorin peroxid 
and some perchloric acid, HC10 4 . 

Chlorin peroxid is a heavy yellow gas, having a peculiar 
odor, a little like that of chlorin and burnt sugar. In 
the cold, it easily condenses to a dark red liquid. In 
both the gaseous and liquid form, it is very explosive. 

It is a powerful oxidizing agent, and many substances 
take fire when brought into contact with it. It is quite 
soluble in water, with which it forms a yellow solution. 

72. The Oxy-acids of Chlorin. There are four oxy- 
acids of chlorin. They are, 

Hypochlorous acid, HCIO, 
Chlorous acid, HC10 2 , 
Chloric acid, HC10 3 , and 
Perchloric acid, HC10 4 . 

These acids are all unstable and can only exist in weak 
aqueous solutions. Only the perchloric acid can be pre- 



64 LECTURE NOTES 

pared in a pure state, and this undergoes decomposition, 
often with violent explosion, when preserved for a few 
days, even in the dark. These acids form salts which 
are quite stable, and some of them are very important 
commercial compounds. All these acids are themselves 
powerful oxidizing agents, which constitutes their prin- 
cipal use. 

7.3. Hypochlorous Acid and its Salts. This acid is 
found only in dilute solutions, and any attempt to con- 
centrate it, results in its decomposition. It is a powerful 
bleaching agent, but since its salts act in a similar way, 
they are more commonly employed for this purpose. 

If chlorin gas is brought into contact with a base it 
forms, in part, a hypochlorite of the metal. The most 
important salt of this acid is bleaching powder, which is 
known commercially as "chlorid of lime.''' 

Bleaching powder is formed, by leading chlorin into 
calcium hydroxid or slaked lime, Ca(OH) 2 , thus : 

2 Ca(OH) 2 + 2 Cl 2 = CaCl 2 + Ca(C10) 2 -f 2 H 2 0. 

Bleaching powder was at first regarded as a mixture of 
calcium chlorid, CaCl 2 , and calcium hypochlorite, Ca(C10) 2 , 
as would seem to be indicated by the equation. It is 
however, believed to be a chemical combination of the two 
salts, and its formula may be written thus : 

Ca < O— CI. 

It is a white porous solid, and has an odor like chlorin. 
It slowly decomposes in the air, or in an aqueous solution, 
setting free hypochlorous acid. It is much used in 
bleaching, and as a disinfectant. 

74. Chlorous Acid and its Salts. Chlorous acid is not 



LECTURE NOTES 65 

known in the free state. A few of its salts, called chlor- 
ites, are known. If we add some soluble base, like 
potassium hydroxid, KOH, to an aqueous solution of 
chlorin peroxid. we obtain a mixture of potassium chlor- 
ite, KC10 2 , and potassium chlorate, KC10 3 . From this 
chlorite, the others may be prepared. 

All chlorites decompose easily and so act as oxidizing 
agents. They all bleach vegetable colors. None of them 
has any commercial value. 

75. Chloric Acid and Its Salts. This acid can be pre- 
pared by decomposing barium chlorate with dilute sulfuric 
acid, thus : 

Ba(C10 3 ) 2 -f H 2 S0 4 = BaS0 4 -f 2 HC10 3 . 

The clear solution of chloric acid may be poured off 
from the insoluble barium sulfate, and concentrated in a 
vacuum over sulfuric acid, until its strength is about 40 
per cent. , when it begins to decompose. 

When chloric acid is allowed to stand for some time 
exposed to the light, it decomposes, forming perchloric 
acid, HC10 4 , liberating oxygen and chlorin. It is, there- 
fore, a powerful oxidizing agent. Many organic sub- 
stances are so rapidly oxidized by it, as to take fire. 

The salts of chloric acid are called chlorates, and the 
most important of these is potassium chlorate, known 
commercially as "chlorate of potash." This is prepared 
by leading chlorin into a solution of potassium hydroxid, 
thus: 

3 Cl 2 + 6 KOH = 5 KC1 -f 3 H 2 -f KC10 3 . 

Potassium chlorate is very much used in the laboratory, 
as an oxidizing agent, and in medicine, for treating cer- 
tain diseases of the throat. 



66 LECTtTRE NOTES 

76. Perchloric Acid. This acid is the most stable of all 
this series of acids, and can be obtained pure, by treating 
potassium perchlorate with concentrated sulfuric acid. 
When pure, it is a colorless liquid, which takes moisture 
rapidly from the air. It can also be obtained in the form 
of crystals, which melt at 15°. It oxidizes organic sub- 
stances so vigorously as often to cause explosion. 

Its salts, which are called perchlorates, are not of 
special importance. 

77. The Constitution of the Oxy-acids of Chlorin. The 

constitution of these compounds has been the subject of a 
great amount of discussion, nor can it be regarded as 
entirely settled. It was at first believed that chlorin was 
always univalent, and, since oxygen is always bivalent, 
the following seemed to be the only constitutional for- 
mulas for these acids: 

Hypochlorous acid, H — O — CI. 
Chlorous acid, H— O— O— CI. 
Chloric acid, H— O— O— O— CI. 
Perchloric acid, H— O— O— O— O— CI. 

In the case of other compounds which contain oxygen 
atoms combined together in this way, we find that the 
compounds become more unstable as the number of 
oxygen atoms increases. In the case of these acids, per- 
chloric acid is the most stable of the four, and the only 
one which can be obtained pure. 

This is best explained by the theory, which is more 
generally accepted at the present time, that the valence 
of chlorin is variable, being one, three, five and seven in 



LECTURE NOTES 67 

the four acids. The following represents their constitu- 
tional formulas according to this theory : 

H— O— CI. O 

H— O— CI ^ ~ H— O— C1=0. 

^ II 

H— 0-Cl=0. O 

The evidence seems to favor this view of the variable 
valence of chlorin, and, furthermore, in chlorin peroxid, 
C10 2 , the valence of chlorin is regarded as being four. 
These views are strengthened by the study of the corre- 
sponding compounds of the other halogen elements. 

Bromin. Symbol, Br. Atomic Weight, 80. 

78. History and Occurrence. Bromin was first discov- 
ered by Balard, in 1826. It was obtained from the liquor 
called bittern, which remains after the sodium chlorid has 
crystallized out from concentrated sea-water. It never 
occurs free in nature on account of its strong affinities. 
It is usually found in connection with compounds of chlorin, 
occurring as bromids of sodium, potassium or magnesium, 
in sea-water, in mineral springs, and in salt deposits, 
such as the one at Stassfurt, in Germany. 

79. Preparation. Bromin can be prepared from bro- 
mids by the same method employed for the preparation 
of chlorin, thus : 

2 NaBr + Mn0 2 -f 2 H 2 S0 4 = 

MnS0 4 -f Na 2 S0 4 -f 2 H 2 + Br 2 . 

It may also be prepared from the bromids by treating 
them with free chlorin, when the bromin is replaced by 
the chlorin, thus : 

2 NaBr + CL = 2 NaCl + Br.. 



68 LECTURE NOTES 

It can be prepared in other ways, analogous to those 
employed for making chlorin. 

80. Properties. Bromin is a heavy, red liquid, which 
is so dark as to be opaque except in very thin layers. It 
is the only element, except mercury, which is a liquid at 
the ordinary temperature. Its specific gravity is 3.18 ; 
it evaporates rapidly in the air at the ordinary tempera- 
ture, and boils at 63°. It is somewhat soluble in water, 
100 parts of which dissolve a little more than three parts 
of bromin. This solution, known as bromin water, is much 
used as an oxidizing agent. It has a very disagreeable 
odor, and its vapor attacks the eyes. When breathed, 
even in very dilute form, it causes great irritation to the 
throat and lungs. It is an irritant poison, and causes 
corrosive sores if dropped on the skin. 

In its chemical properties bromin very closely resem- 
bles chlorin, but its affinities are weaker. It combines 
with most of the elements directly and with some of them, 
particularly potassium, with explosive violence. It may 
be used for bleaching, the same as chlorin. 

COMPOUND OF BROMIN WITH HYDROGEN 

Hydrogen Bromid, or Hydrobromic Acid. Molecular 
Formula, HBr. Molecular Weight, 8i. Density, 40.5. 

81. Preparation and Properties. Hydrogen and bromin 
can be made to combine directly by heat, but they do not 
combine in the sunlight, as is the case with hydrogen and 
chlorin. 

The preparation of hydrobromic acid from its salts, the 
bromids, is complicated by the fact, that most of the acids 
which decompose the bromids will partly decompose the 



LECTURE NOTES 69 

hydrobromic acid which is formed, giving free bromin. 
(Compare hydrochloric acid, 68.) Phosphoric acid could 
be used, but sulfuric acid, which is generally used for 
such processes, must be so diluted, that the results ob- 
tained are not very satisfactory. The best way to prepare 
hydrobromic acid, is to bring bromin and phosphorus 
together in the presence of water. Amorphous phos- 
phorus is placed in water in a suitable vessel, and the 
bromin allowed to drop into it. The action which is at 
first quite violent, gives hydrobromic and phosphoric 
acids, thus : 

p _j_ 5 Br + 4 H 2 = 5 HBr + H 3 P0 4 . 

Hydrobromic acid is a colorless gas. It fumes strongly 
in the air, and is very soluble in water, with which it 
forms a solution which resembles aqueous hydrochloric 
acid. If the solution is saturated at 0°, it has a specific- 
gravity of 1.78 ; but by heating this solution, hydrobro- 
mic acid gas passes off until the specific gravity reduces 
to 1.45, and the liquid contains 48 per cent, of the acid. 
At this strength, the acid may be distilled without further 
change, boiling at 126°. 

82. The Oxids and Oxy-acids of Bromin. No com- 
pound of bromin with oxygen alone, has ever been ob- 
tained. 

There are two oxy-acids known. These correspond to 
two of the oxy-acids of chlorin, and they are named' ac- 
cording to their analogous compounds. They are, 

Hypobromous acid, HBrO. and 
Bromic acid. HBr0 3 . 

They are both quite unstable compounds and can only 
be obtained in an aqueous solution. Their salts are much 



70 LECTURE NOTES 

more stable, and correspond in every way to their analo- 
gous chlorin compounds. 

Iodin. Symbol, I. Atomic Weight, 126.8. 

83. History and Occurrence. Iodin was discovered in 
1812 by Courtois. Before the present methods for mak- 
ing sodium carbonate were discovered, this important salt 
was obtained from sea-plants. These were dried and 
burned, and the salt extracted from the ash by leaching. 
[This ash was known in Scotland as kelp and in Normandy 
as octree.'] It was while examining the mother-liquor, 
which remained after the salts had crystallized out, that 
iodin was discovered. 

Iodin does not occur free in nature, but, in small quan- 
tities, it is quite widely distributed in both organic and 
inorganic compounds. It is found in sea-water, in min- 
eral springs, and especially in plants which grow in the 
deep sea, and in certain sea animals such as sponges and 
oysters. It is also found in the crude Chili saltpeter or 
ealiche, and it is from this source that most of the iodin 
is now obtained. 

84. Preparation. To prepare iodin from sea- weed, the 
mother-liquor, which remains after the soda has crystal- 
lized out, is treated with an excess of sulfuric acid. This 
decomposes most of the compounds and changes the iodin 
present, into hydriodic acid. The liquor is then trans- 
ferred to proper receptacles where it is heated. Manga- 
nese dioxid is added from time to time, when the iodin is 
set free and distils over. The final* action is as follows : 

2 HI + H 2 S0 4 + Mn0 2 = MnS0 4 + 2 H 2 + I 2 . 

After the iodin has all been driven off, some bromin is 
obtained from the residue. 



LECTURE NOTES 71 

In the Chili saltpeter, the iodin exists in the form of 
sodium iodate, NaIO g . This remains in the mother-liquor 
from which it is precipitated by sodium bisulfite, thus : 

2 NaI0 3 + 5 NaHS0 3 =3 NaHS0 4 + 2 Na s S0 4 -f H 2 -f I 2# 

Iodin may be prepared in the laboratory by a method 
which is analogous to that employed for chlorin and bro- 
min, thus : 

2 KI + 2 H 2 SO + + Mn0 2 = 

MnS0 4 + K,S0 4 + 2 H 2 + I 2 . 



85. Properties. Iodin is a dark gray crystalline solid. 
It has a sharp sour taste, and an odor like very weak 
chlorin. It is quite heavy, its specific gravity being 4.95. 
It melts at 114°, and boils at 184°, giving a vapor which, 
when pure, has a splendid deep blue color, but which is 
violet when mixed with air. This is the heaviest vapor 
known, its specific gravity being 8.72. 

Iodin is only slightly soluble in water but is easily 
soluble in an aqueous solution of potassium iodid. It also 
dissolves in alcohol, ether, chloroform, carbon disulfid, 
and many liquid hydrocarbons, imparting to them a fine 
violet, or deep red, color. 

Iodin closely resembles chlorin and bromin in its chem- 
ical properties, but its chemism is less, and so in its 
compounds with the metals it is easily replaced by either 
of the other elements, thus, for example : 

2 KI + C1 2 = 2 KC1 + I 2 . 

Heat decomposes most of its compounds setting free the 
iodin, which may be recognized by the color of its vapor. 

The most characteristic property of iodin is the forma- 
tion of a splendid blue color with starch paste. The blue 



72 LECTURE NOTES 

color disappears on heating the solution, but reappears 
on cooling. 

Iodin stains the skin dark-brown, and is used in medi- 
cine as a local-irritant. Its compounds are also used in 
medicine and in photography. 

COMPOUND OF IODIN WITH HYDROGEN 

Hydrogen Iodid, or Hydriodic Acid. Molecular Formula, 
HI. Molecular Weight, 127.8. Density, 63.9. 

86. Preparation and Properties. Hydrogen and iodin 
do not combine directly by heat alone. They can be 
made to combine, however, by passing them over finely 
divided platinum, heated to redness. 

Hydriodic acid can best be prepared by a method simi- 
lar to that which was employed for the preparation of 
hydrobromic acid, that is, by allowing iodin and phos- 
. phorus to act upon each other in the presence of water, 
thus : 

P _|_ 5 I -f 4 H 2 = 5 HI + H 3 P0 4 . 

Hydriodic acid is a colorless gas which fumes strongly 
in the air. It is very soluble in water and the solution 
very closely resembles aqueous hydrochloric and hydro- 
bromic acids. If the solution is saturated at 0°, it has a 
specific gravity of 1.99 ; but if the solution is heated, 
hydriodic acid gas passes off until the specific gravity is 
reduced to 1.67, and the liquid contains 58 per cent, of 
the acid. This solution can be distilled without further 
change, boiling at 127°. This solution oxidizes in the 
presence of the air, iodin being liberated. 

Hydriodic acid gas is easily decomposed by heat into 
hydrogen and iodin. It is a strong reducing agent, and 



LECTURE NOTES 73 

is much used for this, and other purposes, in organic 
chemistry. 

87. The Oxids and Oxy-acids of lodin. Iodin forms 
only one compound with oxygen, viz. : iodin pentoxid, 

iA- 

This is a white crystalline solid, and quite stable. 
There are two oxy-acids known. They are : 

Iodic acid, HI0 3 , and 
Periodic acid, HI0 4 . 

Iodic acid can be prepared from barium iodate, Ba(IO s ) 2 , 
by the action of dilute sulfuric acid. Also by the oxida- 
tion of iodin by means of concentrated nitric acid, thus : 

3 I 2 + 10 HN0 3 = 6 HI0 3 -f 10 NO + 2 H 2 0. 

Iodic acid is a white crystalline solid, which is easily 
soluble in water. It is reduced by sulfur dioxid, or 
hydrogen sulfid, with separation of free iodin. Its con- 
stitution is similar to that of chloric acid. (Compare 77.) 
Its salts are called iodates. 

The normal periodic acid, HIO + , is not known, but a 
hydrated acid, HI0 4 -j- 2 H 2 0, or, H 5 IO e , is known. It is a 
white crystalline solid and at 140° decomposes into iodin 
pentoxid, water, and oxygen. 

It forms a remarkable series of salts called periodates. 
which may be regarded as derived from a series of hypo- 
thetical acids, which are themselves derived from the 
hypothetical compound, (HO) I, by the loss of varying 
amounts of water. 

The principal value of this series of salts, is in estab- 
lishing the constitution of periodic acid, which is similar 
to perchloric acid. (Compare 77.) 



74 LECTURE NOTES 

Fluorin. Symbol, F. Atomic Weight, 19. 

88. History and Occurrence. It has long been known 
that if the mineral fluorite is acted upon by sulfuric acid, 
a compound is obtained which will etch glass. All 
attempts to isolate fluorin failed until 1886, when Moissan 
succeeded in obtaining the pure element, and in studying 
its properties. 

Fluorin occurs, combined with calcium as calcium fluo- 
rid, CaF 2 , in the mineral fluorite, or fluorspar ; combined 
as a double fluorid of sodium and aluminum, AlF 3 ,3NaF, 
in the mineral cryolite ; in small quantities in many other 
mineral substances ; in minute traces in the bones and 
teeth of animals, in blood and in milk. 

89. Preparation and Properties. The preparation of 
fluorin was one of the most difficult problems of modern 
chemistry. The following is the method employed by 
Moissan for this purpose. 

A solution of acid potassium fluorid, HKF 2 , in pure 
anhydrous hydrofluoric acid, was placed in a U-shaped 
tube of iridioplatinum, in which were placed electrodes 
of the same alloy. The open ends of the tube were closed 
by stoppers of fluorite and the joints made air tight by 
lead washers. Two delivery tubes of platinum served to 
conduct away the gases formed. The whole was then 
placed in methyl chlorid, a liquid which boils at — 23°, 
and then a powerful electric current passed through the 
solution. By this means the hydrofluoric acid was decom- 
posed, hydrogen being liberated at the negative pole, and 
fluorin at the positive pole. 

Fluorin is a light yellow gas, and has an odor something 
like chlorin. It is the most intensely active element 



LECTURE NOTES 75 

known. It cannot be collected by any of the ordinary 
methods since it attacks glass, combines with mercury 
instantly, and decomposes water. It combines with 
hydrogen with explosive violence, even in the dark and 
at a very low temperature. It does not combine with 
oxygen in any form, being the only element which does 
not do so. 

COMPOUND OF FLUORIN WITH HYDROGEN 

Hydrogen Fluorid, or Hydrofluoric Acid. Molecular For- 
mula, HF. Molecular Weight, 20. Density, 10. 

90. Preparation and Properties. Hydrofluoric acid may 
be prepared by the action of sulfuric acid upon any fluorid. 
We generally use calcium fluorid, and the action is repre- 
sented by the following equation : 

CaF, 4- H SO, = CaSO, + 2 HF. 

2 12 4 4 1 

Pure hydrofluoric acid is a colorless and very volatile 
liquid, which fumes in the air, and boils at 19°. It is 
easily soluble in water, and is usually employed in this 
form. 

The most remarkable property of hydrofluoric acid is its 
etching of glass. Glass is composed, principally, of sili- 
con dioxid, and it is the affinity of fluorin for the silicon 
which causes the etching. The two elements combine to 
form silicon tetrafluorid, SiF 4 , thus : 

4HF + Si0 2 = SiF 4 -J- 2 H 2 0. 

The glass to be etched, is first coated with wax, and 
the design cut through the wax with a sharp instrument. 
It is then exposed to the action of the acid, either in the 
form of gas, or in an aqueous solution. 

Hydrofluoric acid is kept in bottles made either of cere- 



76 LECTURE NOTES 

sine, or hard rubber, or, better still, of platinum. The 
anhydrous acid is a powerful caustic, destroying the skin, 
and causing painful sores which are difficult to heal. 

Just above the boiling point, the density of the pure 
acid corresponds to the molecule H 2 F 2 . At this temper- 
ature the fluorin atom is therefore trivalent, and the 
constitutional formula is H— Fr=F — H. 

The salts of hydrofluoric acid are called fluorids. 

91. The Sulfur Group. The three elements, sulfur, 
selenium, and tellurium, form what is known as the sulfur 
group. They differ considerably in their physical prop- 
erties, but chemically are much alike, and form compounds 
in every way analogous. Sulfur is a very common element 
while selenium and tellurium are quite rare. 

In the periodic arrangement of the elements, oxygen 
occurs as a member of this group, and in most of its 
chemical properties it very closely resembles sulfur. 

In compounds with the metals and with hydrogen, these 
elements are bivalent and very much resemble oxygen. 
They form compounds with oxygen, however, and in them 
exhibit great variation in valence, being tetrads and 
hexads in some of them. They all exhibit a strong nega- 
tive character. This decreases as their atomic weights 
increase, sulfur having the strongest affinities and tellu- 
rium the weakest. 

Sulfur. Symbol, S. Atomic Weight, 32. 

92. History and Occurrence. Sulfur has been known 
since the very earliest times. It was very highly regarded 
by the alchemists, and believed by them to be the principle 
of combustibility. 



LECTURE NOTES 77 

Sulfur is a very widely distributed element, and occurs 
both free, and combined with other elements. Free sulfur 
is found in volcanic districts, particularly in Sicily, where 
it is in part formed by the mutual decomposition of certain 
volcanic gases. 

The compounds of sulfur found in nature are very 
numerous, and they occur in all parts of the world. Of 
these, the most important compounds are the sulfids and 
sulfates. The principal sulfids arepyrite, or iron pyrites. 
FeS 2 ; chalcopyrite, CuFeS 2 ; galenite, PbS ; sphalerite, 
ZnS ; and cinnabar, HgS. The most important sulfate is 
gypsum, CaS0 4 . 

Sulfur is an important constituent in many organic 
compounds. It comprises about one per cent, of all albu- 
minous substances. It is found in small quantities in 
wool and in hair, also in mustard and garlic. It is the 
presence of sulfur in these compounds that is the principal 
cause of the disagreeable odors when they decay. 

93. Preparation. • Most of the sulfur of commerce is 
obtained from the native material found in Italy. This 
is usually found mixed with marl, gypsum, celestine, and 
earthy impurities. To free it from these, the early 
method was to place heaps of the impure sulfur in holes 
dug in the ground, and set them on fire. The heat of the 
sulfur burning on the surface, melted the rest of it, 
which ran to the bottom, and was removed fairly free from 
the impurities. By this process only about one- third of 
the sulfur was obtained, the rest being burned for fuel. 
A special kiln is now employed, which is operated on 
a somewhat more economical plan; but the principle is 
the same, the crude sulfur being still the cheapest fuel 
for the purpose. 



78 LECTURE NOTES 

Other methods of extracting the sulfur have been tried, 
but they have not proved successful. 

The crude sulfur, thus obtained, is afterwards purified 
by distillation, the vapor being led into a cold chamber, 
where it condenses, forming the so-called flowers of sul- 
fur. After a time the chamber becomes heated above 
the melting point of sulfur. The liquid sulfur is then 
drawn off, and cast in wooden moulds in the form which 
we call roll sulfur, or brimstone. 

Sulfur may be prepared by heating pyrite, the action 
being as follows : 

3 FeS 2 = Fe s S 4 + S 2 

A very interesting method for preparing it is the way 
in which it is probably formed in nature. This is by 
the action of sulfur dioxid upon hydrogen sulfid, both of 
which gases are found issuing from all active volcanoes. 
This action is as follows : 

SO, -f- 2 H S = 2 HO + 3S. 

2 I 2 2 I 

Another important source of sulfur is the residue or 
waste from the manufacture of sodium carbonate by the 
Leblanc process. This waste material consists principally 
of calcium sulfid and other impurities. This is mixed with 
water, and a stream of carbon dioxid gas passed through 
the mixture, when hydrogen sulfid is produced, thus : 

CaS -f H 2 + C0 2 = CaCO s + H 2 S. 

The hydrogen sulfid is then burned with an insufficient 
supply of air and the sulfur recovered, thus : 

H 2 S + O = H 2 + S. 
94. Properties. Sulfur is a yellow crystalline solid. 



LECTURE NOTES 79 

It has a very slight odor when rubbed. It is exceeding- 
ly brittle, and, if a piece is held in the hand, it is heard 
to crackle. It is insoluble in water but dissolves easily 
in carbon disulfid, and, to some extent, in many organic 
liquids. 

Sulfur melts at 114.5°, forming a thin, amber-colored 
liquid. If it is further heated it soon becomes dark red 
and quite viscid. At about 230° it is almost black, and 
so thick that it can hardly be poured from the vessel. If 
the heating is continued the liquid becomes less viscid but 
the dark color remains. At 448° it boils and forms a 
dark red vapor. If the boiling sulfur is cooled the same 
changes occur in the reverse order until it solidifies. 

At a little above the boiling point, say 500°, its 
density is found to be 96. This would make its molecular 
weight at this temperature 192 ; and as its atomic 
weight is 32, at 500° its molecule must contain six 
atoms. At 1000° its density is 32, so that at 1000° its 
molecule contains but two atoms, as is the case with 
most of the elementary gases. 

When sulfur is heated in the air to about 260°, it takes 
fire and burns with a pale blue flame, forming sulfur 
dioxid, S0 2 , which has a characteristic, suffocating odor. 
In the gaseous form, sulfur supports the combustion of 
many of the metals, thus showing its likeness to oxygen. 

If finely powdered sulfur is moistened with water and 
exposed to the air, it is slowly oxidized, forming sulfuric 
acid. The compounds of sulfur with the metals are called 
sulfids. 

95. The Allotropic Forms of Sulfur. Sulfur is capable 
of existing in at least four allotropic forms, two of which 



80 LECTURE NOTES 

are crystalline and two amorphous. These different forms 
are probably due to changes in the number of atoms in 
the molecule. 

The first crystalline variety is known as rhombic sulfur. 
It is also called octahedral sulfur, and a-sulfur. It is 
the form in which sulfur occurs in nature. It is formed 
when sulfur is allowed to crystallize from a solution, and 
may be obtained, under certain conditions, when melted 
sulfur is allowed to crystallize. It is a lemon-yellow crys- 
talline solid, has a specific gravity of 2.05 and is easily 
soluble in carbon disulfid. 

The second crystalline variety is known as monoclinic 
sulfur. It is also called prismatic sulfur, and (6-sulfur. 
It is formed when melted sulfur is allowed to cool quickly. 
To prepare this variety, sulfur is melted in a crucible and 
then allowed to cool. As soon as a crust is formed on the 
surface, it is broken through, and the portion of sulfur 
which is still liquid, is poured out. The crystallized 
portion is seen to consist of long, transparent, brownish- 
yellow needles. At the ordinary temperature this variety 
is unstable, so that after a day or two the crystals lose 
their transparency, become lemon-yellow in color, and are 
found to have changed to the rhombic variety. Mono- 
clinic sulfur melts at 120°, has a specific gravity of 1.96, 
and is soluble in carbon disulfid. 

The change from one crystalline variety to the other 
takes place at about 100°. If the melted sulfur, which 
ordinarily forms the monoclinic variety, is cooled very 
slowly, it can be kept in the liquid form until the tem- 
perature is only about 90°. If kept at this temperature 
for some time, rhombic crystals will form. 

The first amorphous variety of sulfur is known as pi as- 



LECTURE NOTES 81 

tic sulfur or y -sulfur . It is formed when sulfur, which 
has been heated to 250° or more, is suddenly cooled by 
pouring it into water. It is a dark brown tenacious mass, 
resembling rubber, has a specific gravity of 1.95 and is 
insoluble in carbon disulfid. 

The second amorphous variety is known as milk of sulfur 
or lac sulfur. It is formed when the soluble polysulfids 
are acted upon by hydrochloric acid. It is a milk white, 
very finely divided powder, is quite stable and easily 
soluble in carbon disulfid. It gradually changes to the 
rhombic variety at the ordinary temperature, and does so 
quickly at 100°. There are possibly other varieties of 
sulfur, but those described as such are probably mixtures 
of these varieties. 

COMPOUNDS OF SULFUR WITH HYDROGEN 

Sulfur forms two compounds with hydrogen. These are: 

Hydrogen sulfid, H 2 S, and 
Hydrogen persulfid, H 2 S 2 . 

The latter compound has little but theoretical impor- 
tance ; the former is one of the most important compounds 
known. 

Hydrogen Sulfid. Molecular Formula, H 2 S. Molecular 
Weight, 34- 

96. History and Occurrence. Hydrogen sulfid was dis- 
covered in 1777 by Scheele, who made it by heating sul- 
fur in hydrogen gas. Although many processes were 
known to the alchemists, during which hydrogen sulfid 
must have been formed, it had always been classified 
by them under the general head of sulfurous vapors. 

Hydrogen sulfid occurs in nature in volcanic gases and 



82 LECTURE NOTES 

in certain natural waters, the so-called sulfur springs. 
It is also formed in nature by the decay of organic matter 
containing sulfur. 

97. Preparation. Hydrogen sulfid is easily obtained 
by the action of an acid upon some sulfid. Ferrous sulfid 
and sulfuric or hydrochloric acid are usually employed, 
the action taking place in the cold according to the follow- 
ing equation : 

FeS + H 2 S0 4 = FeS0 4 -f H 2 S. 

Since the ferrous sulfid is never pure, but always con- 
tains some metallic iron, the hydrogen sulfid, if prepared 
by this method, always contains some hydrogen. 

If hydrogen sulfid is required pure, antimony sulfid and 
hydrochloric acid are used, thus : 

Sb 2 S 3 + 6 HC1 == 2 SbCl 3 -f 3 H 2 S. 

Hydrogen sulfid may also be obtained by direct union 
of the elements, by leading hydrogen gas through boiling 
sulfur. 

98. Properties and Uses. Hydrogen sulfid is a color- 
less gas. It has a sweetish taste, and a very disagreeable 
odor, quite similar to that of rotten eggs. It is soluble 
in water, one volume of which dissolves about three 
volumes of the gas, at the ordinary temperature. This 
solution possesses weak acid properties, and is sometimes 
called hydro sulfuric acid. The sulfids are regarded as 
salts of this acid. A solution of this gas decomposes 
slowly on standing, especially in the sunlight, forming 
water, and depositing sulfur. 

It is an inflammable gas, and burns with a pale blue 
flame. It* two volumes of the gas are mixed with three 



LECTURE NOTES 83 

volumes of oxygen, and ignited, they explode with consid- 
erable violence. 

It is an active poison, less than one per cent, in the 
atmosphere being dangerous to breathe. The halogen 
elements decompose it easily, and so dilute chlorin is used 
as an antidote for poisoning with this substance. 

It is one of the most useful reagents in the laboratory. 
It forms insoluble sulfids with a number of the metals, 
and so is used to precipitate the metals from their solu- 
tions. It is also a reducing agent, and is often used for 
this purpose. 

99. The Halogen Compounds of Sulfur. Sulfur forms 
three compounds with chlorin, each by direct union of 
the elements. These are : 

Sulfur monochlorid, S 2 C1 2 , 
Sulfur dichlorid, SC1 2 , and 
Sulfur tetrachlorid, SCI . 

Sulfur monochlorid is formed by leading dry chlorin 
into melted sulfur. It is an amber-colored liquid, has a 
very disagreeable, penetrating odor, and boils at 138°. 
It dissolves sulfur very easily and the solution thus ob- 
tained is largely employed in vulcanizing rubber. 

The dichlorid is formed by leading chlorin into the 
monochlorid at 0°, and the tetrachlorid, by leading chlorin 
into the dichlorid at — 22°. 

Sulfur forms one bromid, S 2 Br 2 , and one iodid, S 2 I 2 , 
each by direct union of the elements. 

No compound of sulfur with fluorin is known to exist. 



84 LECTURE NOTES 

COMPOUNDS OF SULFUR WITH OXYGEN 

Sulfur forms four compounds with oxygen. These are : 

Sulfur dioxid, S0 2 , 

Sulfur trioxid, S0 3 . 

Sulfur sesquioxid, S 8 8 , and 

Sulfur heptoxid, S 2 7 . 
Only the first two are of much importance. 

Sulfur Dioxid. Molecular Formula, S0 2 . Molecular 
Weight, 64. 

100. History and Occurrence. We have already noted 
that sulfur has been known since the earliest times. The 
ancients knew that when sulfur was burned, pungent 
vapors were evolved, and they employed them for fumi- 
gation and bleaching. They also knew that this vapor 
was not the same as sulfuric acid. In 1775, Priestley 
first prepared the pure gas which he called sulfurous acid. 

Sulfur dioxid is found in nature in volcanic gases, and 
in solution in certain volcanic springs. 

101. Preparation. Sulfur dioxid is easily prepared 
synthetically, by burning sulfur in the air or in oxygen. 

For use in the laboratory, it is most conveniently pre- 
pared by the action of certain metals, such as copper, 
mercury, or silver, on hot concentrated sulfuric acid, thus: 

Cu -f 2 H 2 S0 4 = CuS0 4 -f 2 H 2 -j- S0 2 . 

In this action hydrogen is first produced, which, in the 
nascent state, decomposes the hot sulfuric acid. 

If a sulfite is acted upon by dilute sulfuric acid, sulfur- 
dioxid is liberated, thus : 

Na 2 SO s + H a S0 4 = Na 2 S0 4 + H 2 + S0 2 . 



LECTURE NOTES 85 

Sulfur dioxid is made on a large scale for commercial 
purposes by heating charcoal (carbon) with concentrated 
sulfuric acid, thus : 

C + 2 H 2 S0 4 = 2H 2 0+ C0 2 + 2 S0 2 . 

102. Properties. Sulfur dioxid is a colorless gas. It 
has a sharp, sour taste, and a peculiar suffocating odor, 
like a burning sulfur match. It does not burn, neither 
does it support combustion. 

Sulfur dioxid is quite easily soluble in water. At 0° 
one volume of water dissolves about 80 volumes of the 
gas, and at the ordinary temperature (20°) it dissolves 
about 40 volumes. This solution has strongly acid prop- 
erties, and is known as sulfurous acid, the sulfur dioxid 
and water having combined chemically to form the acid, 
thus : 

S0 2 + H 2 = H 2 S0 3 . 

When this solution is boiled, all of the sulfur dioxid is 
expelled. 

Sulfur dioxid is easily condensed to a colorless, mobile 
liquid. This requires a pressure of about three atmos- 
pheres at 15°, or a temperature of — 10° at the ordinary 
atmospheric pressure. If the liquid sulfur dioxid is 
allowed to evaporate, a great amount of heat is rendered 
latent, thus producing great cold. If a stream of dry air 
is driven through the liquid, a temperature of — 50° is 
easily obtained. 

Sulfur dioxid, either as a gas, or in solution, possesses 
powerful bleaching properties, and it is used in the arts 
for bleaching silk, wool, and straw goods, which are injured 
by chlorin. The bleaching action of sulfur dioxid depends 



86 LECTURE NOTES 

upon the fact, that in the presence of water it is oxidized 
and hydrogen is liberated, thus : 

S0 2 + 2H 2 = H 2 S0 4 + H 2 . 

The hydrogen, thus liberated, reduces the coloring 
matter forming a colorless compound. The bleaching 
action of sulfur dioxid is, therefore, one of reduction, 
while the bleaching action of chlorin is one of oxidation. 

Sulfur dioxid possesses powerful antiseptic properties, 
and so is much used as a disinfectant. It is used in the 
laboratory as a reducing agent. 

Sulfur Trioxid. Molecular Formula, S0 3 Molecular 
Weight, 80. 

103. Preparation and Properties. Sulfur trioxid, called 
also sulfuric anhydrid, can be formed synthetically by 
leading a mixture of sulfur dioxid and oxygen, over heated 
platinum-sponge, or platinized asbestus. 

It is most conveniently prepared by heating pyrosulfu- 
ric acid. This compound is easily decomposed by heat, 
forming sulfuric acid and sulfur trioxid, thus : 

H s S s O, = H,SO t + SO„. 

Pure sulfur trioxid at the ordinary temperature, is a 
colorless liquid, which boils at 46°. If cooled, it forms 
crystals which melt at 14.8°. Itis very difficult to obtain 
this substance pure, and free from water, and if a very 
small amount of water is present, it forms a mass of white 
silky needles, which do not melt until about 50°. 

Sulfur trioxid has great affinity for water, with which 
it forms sulfuric acid. If it is thrown into water, it dis- 
solves with a hissing sound, and gives off a great amount 
of heat. It gives dense white fumes when exposed to the 



LECTURE NOTES 87 

air, owing to the absorption of moisture. At a red heat 
it is decomposed into sulfur dioxid and oxygen. 

COMPOUNDS OF SULFUR WITH OXYGEN AND 
HYDROGEN 

' 104. The Oxy-acids of Sulfur. There are two very 
important oxy-acids of sulfur. These are : 

Sulfurous acid, H 2 SO s . and 
Sulfuric acid. H 2 S0 4 . 

In addition to these, there are no less than nine others. 
These are : 

Hyposulfurous acid, HS0 2 , or H 2 S 2 4 , 
Pyrosulf uric acid, H 2 S 2 7 , 
Persulfuric acid. HS0 4 , or H 2 S 2 O g . 
Thiosulfuric acid, H 2 S 2 3 , 
Dithionic acid, H 2 S 2 6 , 
Trithionic acid, H 2 S 3 6 . 
Tetrathionic acid, H 2 S 4 6 , 
Pentathionic acid, H 2 S O fi , and 



Hexathionic acid, H S O 



The second of these acids, pyrosulfuric acid, is closely 
related to sulfuric acid. The others are very unstable, 
most of them being unknown in the free state. They all 
form salts which are quite stable, some of them being of 
considerable importance. 

105. Sulfurous Acid, H 2 S0 3 . This is a weak, unstable 
acid, which is only known in an aqueous solution. It is 
formed when sulfur dioxid is dissolved in water. (Com- 
pare 102.) This solution smells like sulfur dioxid and 
gives a strong acid reaction. It oxidizes slowly in the 
air forming sulfuric acid. It is a dibasic acid and its 



88 LECTURE NOTES 

salts, which are called sulfites, are quite important. If a 
solution of a base is saturated with sulfur dioxid an acid 
sulfite is formed, thus : 

NaOH -f S0 2 = NaHS0 3 . 

If this compound is mixed with more of the base, the 
normal salt is formed, thus ; 

NaHS0 3 + NaOH = Na 2 S0 3 -f- H 2 0. 

The normal sulfites are crystalline solids and have no 
odor. If they are acted upon by dilute acids they are 
easily decomposed, sulfur dioxid being liberated, thus : 

Na 2 S0 3 + H 2 S0 4 == Na 2 S0 4 -f H 2 + S0 2 . 

Sulfurous acid is used in the laboratory for the same 
purposes as sulfur dioxid. 

Sulfuric Acid. Molecular Formula, H 2 S0 4 . Molecular 
Weight, 98. 

106. History and Occurrence. Sulfuric acid is un- 
doubtedly the most important and useful acid known. It 
was probably known to the Arabian alchemists in an im- 
pure state, as early as the eighth century. A process 
for preparing it from ferrous sulfate was first described 
by Basil Valentine, in the 15th century. This method 
was the only one employed until the early part of the 
18th century, when the present process of manufacture 
was discovered. 

It was at first made in large glass globes and known 
as sulfuric acid "made by the bell,'' in distinction from 
that made from ferrous sulfate. The production of the 
acid in lead chambers was first proposed by Dr. Roebuck, 
and such chambers were first erected in Birmingham in 
1746. 



LECTURE NOTES 89 

Sulfuric acid was called by the alchemists and early 
chemists, oil of vitriol, because made from ferrous sul- 
fate, or green vitriol, all sufates being called by them 
vitriols. 

Sulfuric acid occurs free in the waters of certain rivers, 
and mineral springs. It is found also in the saliva of 
certain mollusks. Its salts, the sulfates, are found in 
all parts of the world, forming the minerals gypsum, 
CaS0 4 , barite, BaS0 4 , and many others. 

107. Preparation. Sulfuric acid can be formed syn- 
thetically, by adding water to sulfur trioxid. It can also 
be formed by the oxidation of sulfur, or sulfur dioxid, by 
means of nitric acid. The latter is the method employed 
for its manufacture on a large scale. 

Sulfur dioxid is formed, either by burning sulfur in the 
air, or by roasting certain sulfur ores in a suitable furnace. 
Pyrite is generally used for this purpose. Nitric acid 
and steam are then brought into contact with the sulfur 
dioxid and air, when the following action takes place : 

3 SO a -f 2 HNO s + 2 H 2 = 3 H 2 S0 4 + 2 NO. 

The nitric oxid, NO, formed in this action, is then acted 
upon by the air forming nitric peroxid, thus : 

2 NO + 2 = 2 N0 2 . 

The nitric peroxid is then capable of oxidizing more 
sulfur dioxid, and does so as follows : 

S0 2 + N0 2 + H 2 = H 2 S0 4 + NO. 

Sulfur dioxid does not oxidize rapidly in the presence 
of air and water alone, but does so in the presence of the 
higher oxids of nitrogen. The nitric oxid acts thus as a 



90 LECTURE NOTES 

carrier, between the oxygen of the air and the sulfur 
dioxid. 

108. Manufacture. The manufacture of sulfuric acid is 
such a very important industry that it should be thoroughly 
understood by every chemist. All of the sulfuric acid of 
commerce is now made by the lead chamber process, the 
principles involved being those already stated in the pre- 
vious article. 

This process, which is a continuous one, consists in 
bringing the different substances together in a series of 
large chambers lined with lead. The sulfur dioxid is 
usually obtained by roasting pyrite, and the nitric acid 
from sodium nitrate, known commercially as Chili salt- 
peter. The air, together with the hot sulfur dioxid is 
drawn along by the draft from a large chimney, and made 
to pass through what is called a Glover tower, or denitra- 
ting tower, where the heat decomposes certain products 
of a later stage of the operation. From this tower they 
pass into the chambers where the action takes place. 
Steam is also conducted into different parts of the 
chambers. 

The nitric oxid, which is formed in the first operation, is 
oxidized in the presence of air, forming nitric peroxid. 
This can oxidize more sulfur dioxid, forming again nitric 
oxid, when the process is repeated. Theoretically, only 
a very small quantity of nitric acid is necessary to oxid- 
ize a large amount of sulfur dioxid. This would be the 
fact were it not for the large amount of nitrogen in the 
air, which so dilutes the gases that they would soon be- 
come inactive. 

To avoid this, and also to avoid wasting the oxids of 



LECTURE NOTES 91 

nitrogen, which would be carried off through the chimney, 
the gases, after leaving the chambers, pass into what is 
called a Gay-Lussac tower. This is a tall tower, lined with 
lead, and filled with coke, over which strong sulfuric acid 
is made to run. The oxids of nitrogen are absorbed by 
the sulfuric acid, with formation of nitrosulfonic acid, 
N0 2 HS0 3 , known as chamber acid crystals. This collects 
at the bottom of the tower, together with the excess of 
sulfuric acid, while the nitrogen passes on and escapes 
through the chimney. 

The mixed acids are then pumped up into a tank at the 
top of the Glover tower. This tower is also lined with 
lead and filled with flint stones, over which the acids are 
allowed to run slowly. Here they meet the hot gases 
from the sulfur furnaces and are denitrated, thus : 

2 NO,HSO, + SO, + 2 HO = 3 H,S(X + 2 NO. 

The liberated nitric oxid, together with the excess of 
other gases, pass into the lead chambers to go through 
the same process again, the sulfuric acid passing into a 
proper receptacle at the bottom of the tower. 

These two towers are in no way essential to the forma- 
tion of the sulfuric acid, being added for the purpose of 
saving the oxids of nitrogen, which would otherwise 
become waste material. 

The dilute acid, as it comes from the lead chambers, 
contains only about 60 per cent, of pure acid. This is 
concentrated in lead pans until it acquires a strength of 
about 78 per cent., at which strength it begins to attack 
the lead. The concentration is then continued in vessels 
of platinum or glass, until it reaches about 98 per cent. , 



92 LECTURE NOTES 

which is the strength usually found in the concentrated 
acid. 

109. Properties and Uses. Sulfuric acid, when pure, 
is a colorless, odorless, oily liquid. It has a specific gravity 
of 1.84. When the pure acid is heated, it begins to fume, 
and partially decomposes, at 30°, giving off sulfur trioxid. 
This slight decomposition continues, with the increase in 
temperature, until at 338° it boils without further change. 
At a little above 400° it is entirely decomposed into sulfur 
trioxid and water, as shown by the density of its vapor, 
which is found to be only 24.5 at this temperature. These 
two substances combine again on cooling. 

The concentrated acid has a very strong attraction for 
water, the union of the two being attended with the evo- 
lution of a great amount of heat and a contraction of 
volume. In mixing the two, the acid should always be 
poured into the water, and never the reverse. 

Concentrated sulfuric acid does not act upon many of 
the metals when cold, but does in some cases when heated, 
in which case sulfur dioxid is formed. (Compare 101.) 
The dilute acid acts upon the more positive metals, with 
evolution of hydrogen, and the formation of sulfates. 
Many organic substances are decomposed by sulfuric acid, 
which takes from them the elements of water, setting free 
carbon. This explains the charring of wood, paper, sugar, 
etc., when acted upon by this acid. 

Nearly all salts are decomposed by sulfuric acid, giving 
the acids from which the salts were derived. If the acid 
which is produced is unstable, or does not exist free, it is 
further decomposed, giving some characteristic gas. 

The attraction which sulfuric acid has for water is util- 



LECTURE NOTES 93 

ized in the laboratory for the purpose of desiccating, or 
drying gases. 

Next to water, sulfuric acid is the most important com- 
pound known to the chemist, and there is hardly an 
article manufactured, or a chemical process known, which 
does not, directly or indirectly, employ sulfuric acid. 

110. Pyrosulfuric Acid, H 2 S 2 T . This is also known as 
fuming or Nordhausen sulfuric acid, because it was first 
manufactured at Nordhausen, in the Hartz Mts. It may 
be regarded as a solution of sulfur trioxid in sulfuric acid. 
It is also called disidfuric acid, since it may be regarded 
as having been derived from two molecules of sulfuric acid, 
by the removal of one molecule of water. 

Pyrosulfuric acid is best prepared, by first roasting 
ferrous sulfate in the air, which oxidizes it to a basic ferric 
sulfate. This is then heated in clay retorts, giving sulfur 
trioxid, which dissolves in sulfuric acid, forming this acid. 

It is a thick, brown, oily liquid, which, when cooled, 
forms crystals, melting at 35°. It fumes strongly in the 
air, owing to the escape of the volatile sulfur trioxid. It 
is a dibasic acid, forming a well defined class of .salts, and 
is used for dissolving indigo, and in the manufacture of 
alizarin and other organic dyes. 

111. Thiosulfuric Acid, H 2 S 2 3 , and its Salts. This 
compound is better known under its old name of ' ' hyposul- 
furous acid," a name by which we now designate another 
of the oxy-acids of sulfur, viz. : H 2 S,0 4 . Thiosulfuric acid 
is not known in the free state, since even in a very dilute 
form, it soon decomposes, forming sulfur and sulfurous 
acid. 

The salts of this acid, the thiosulfates, are also known 



94 LECTURE NOTES 

commercially as " hyposulfites." The most important 
of these is the sodium salt. This is formed when a solu- 
tion of sodium sulfite is boiled with sulfur, thus : 

Na 2 S0 3 + S = Na s S,O s . 

The soluble thiosulfates show a strong tendency to 
form soluble double salts, and it is this action that gives 
to sodium thiosulfate its principal commercial value. It 
is used for "fixing" the prints in photography. The 
photographic paper is impregnated with silver chlorid, 
which, when exposed to the light, blackens, and becomes 
insoluble in sodium thiosulfate, that which has not been 
exposed, being readily soluble, thus : 

Na 2 S 2 3 + AgCl = AgNaS 2 3 -f NaCl. 

The sodium thiosulfate is known commercially as u hy- 
posulfite of soda," and still more commonly as " hypo." 

Selenium. Symbol, Se. Atomic Weight, 79.2. 

112. History and Properties. Selenium was discovered 
by Berzelius is 1817, in the deposits from the lead cham- 
bers of a sulfuric acid factory in Sweden. 

Selenium is quite a rare element. It is somewhat 
widely distributed, but occurs in very small quantities. 
It is occasionally found free, but generally combined, as 
a selenid of copper, lead, or silver. It is mostly obtained 
from the sulfuric acid residues. 

Selenium exists in several allotropic forms, which may 
be distinguished by their solubility, or insolubility in 
carbon disulfid. Soluble selenium is a reddish-brown or 
black crystalline powder, which has a specific gravity of 
4.3 to 4.5. It crystallizes from a solution in carbon di- 
sulfid in dark-red crystals. It has no definite melting 



LECTURE NOTES 9-5 

point, but softens gradually on heating. The insoluble, 
or metallic selenium is a dark-brown or black crystalline 
solid, having a specific gravity of 4. 5 to 4. 8. It melts at 
217°, and boils at 680°. Metallic selenium conducts 
electricity, and in this connection exhibits some remark- 
able properties, its conductivity being about twice as 
great in the light as in the dark. 

Selenium burns with a blue flame, forming the oxid, Se0 2 , 
which has an exceedingly disagreeable odor, resembling 
that of rotten horse-radish. It unites with the metals, 
directly, to form selenids, and forms a number of acids 
and other compounds which are analogous to those of sulfur. 

Tellurium. Symbol, Te. Atomic Weight, 125. (?) 

113. History and Properties. Tellurium was discovered 
by Klaproth, in 1798, in a gold ore from Transylvania. 
It was at first supposed to be a metal, but its compounds 
so closely resembled those of sulfur and selenium, as to 
place it in the sulfur group. 

Tellurium is a very rare element. It occurs occasionally 
free, but generally as a telluridof gold, silver or bismuth. 
It is a tin-white solid, quite brittle, and a good conductor 
of heat and electricity. Its specific gravity is 6.25. It 
is quite metallic in its physical properties, and exists both 
in the crystalline and amorphous form, both forms being 
insoluble in carbon disulfid. 

It melts at 452°, and, when heated in the air, burns 
with a blue flame forming the oxid, Te0 2 . It forms acids 
and other compounds which are analogous to those of 
sulfur and selenium. Tellurium is too rare an element to 
be of much importance. 



96 LECTURE NOTES 

114. The Nitrogen Group. The elements of this group 
are nitrogen, phosphorus, arsenic, antimony, bismuth, 
and some rare elements. They differ very much physi- 
cally, and to a considerable extent, chemically. Nitrogen 
is a gas, and is one of the least active of all the elements. 
All the others are solids, and phosphorus is one of the 
most active of elements, while antimony and bismuth are 
usually classed among the metals, the latter, especially, 
exhibiting marked metallic properties. 

They aliform compounds with hydrogen, except bismuth, 
and in these compounds are trivalent ; while with oxygen, 
most of them exhibit also a higher valence, being often 
pentads. The compounds of the first three elements show 
much analogy, although their affinities vary considerably. 
Nitrogen is the most common, as well as important ele- 
ment, and may be taken as the type of the group. 

Nitrogen. Symbol, N. Atomic Weight, 14. 

115. History and Occurrence. Nitrogen was discovered 
in 1772 by Dr. Rutherford. He showed that when animals 
breathe in a closed vessel, the air not only becomes laden 
with impurities from the respiration, but the residual air, 
freed from these impurities, will not support combustion 
or respiration. Its relations to oxygen in the atmosphere 
were not known for some time. Lavoisier gave it the 
name azote, because of the above mentioned properties. 
The name nitrogen was afterwards given to it by Chaptal. 

Nitrogen occurs free in the atmosphere, of which it 
constitutes about 78 per cent., by volume. The remain- 
ing portion of the atmosphere consists principally of 
oxygen, with which the nitrogen is mixed, but not chem- 
cally combined. 



LECTURE NOTES 97 

It occurs combined, in ammonium compounds, and in 
nitrates, of which the most important are those of potas- 
sium, sodium, and calcium. It also forms a very essential 
part of many animal and vegetable substances. 

116. Preparation. Nitrogen is most conveniently pre- 
pared by removing oxygen from the atmosphere. This 
may be done in many ways, but is usually accomplished 
by burning phosphorus in a bell-jar over water, or by 
conducting the air over red hot copper. These elements 
combine with the oxygen, leaving nitrogen. 

It may be obtained by heating ammonium nitrite, the 
reaction being as follows : 

NH 4 N0 2 = 2 H 2 + N 2 . 

It may also be obtained by heating ammonium bichro- 
mate, chromium oxid and water being formed, thus : 

(NH 4 ) 2 Cr 8 0,= Cr 2 O s + 4 H 2 + N,. 

A mixture of potassium bichromate and ammonium 
chlorid give a similar result and the cost is much less. 

Other ammonium compounds, especially in the presence 
of oxidizing agents, will, when heated, give nitrogen. 

117. Properties. Pure nitrogen is a colorless, odorless, 
and tasteless gas. It is somewhat lighter than air. its 
specific gravity being 0.97. At a very low temperature, 
and under very strong pressure, it can be condensed to a 
colorless liquid. It is only very slightly soluble in water. 

Nitrogen does not burn, nor will it support the com- 
bustion of anything. It is distinguished by its inactive 
properties, and can be made to combine directly with 
but very few non-metals. At a high temperature, it will 



98 LECTURE NOTES 

combine directly with magnesium, aluminum, and a few 
other metals, forming nitrids. Indirectly, it combines 
with a large number of the elements, and these com- 
pounds often possess remarkably active and character- 
istic properties. 

Many compounds of nitrogen are unstable at a high 
temperature, and it is a constituent of most high ex- 
plosives. 

Nitrogen is not poisonous, although death would ensue 
in an atmosphere of nitrogen, owing to suffocation. Some 
of its compounds, however, are among the most active 
poisons known, and all of the exceedingly poisonous alka- 
loids contain this element. 

The Atmosphere. A Mixture of Nitrogen and Oxygen. 

Specific Gravity, i. Density, 14.39. 

118. Occurrence and Properties. The atmosphere is 
the name given to the invisible, gaseous envelope sur- 
rounding the earth, and at the bottom of which, we live. 

We are aware of its presence by its resistance, motion, 
pressure, and weight. Its resistance is noticed when a 
body passes rapidly through it, or, when it is itself in 
rapid motion, that is, when there is what we call a wind. 

Its pressure, being equal in all directions, is not noticed 
until removed in some direction. The pressure is meas- 
ured by the barometer, and is equal to 1033.3 grams on 
one square centimeter, or about 15 pounds on one square 
inch, of surface. 

Its weight may be shown by weighing a strong flask 
fitted with a stopcock, and after exhausting the air, re- 
weighing the flask when it will be found to have lost 



LECTURE NOTES 99 

weight. Careful determinations show that one liter of 
air, at the normal temperature and pressure, weighs 
1.2935 grams. 

The height to which the atmosphere extends is not 
accurately known, although observation of the time during 
which the twilight curve extends to the zenith, shows that 
this is probably not less than 50 miles, and, possibly, in 
a very attenuated form, it may extend to a much greater 
distance. The reason for this uncertainty is that the 
density decreases very rapidly, but not uniformly, as we 
ascend. If the atmosphere were of the same density 
throughout, its height would be 5.2 miles. 

Since oxygen is necessary to all animal life, the atmos- 
phere, which is the source of supply for this important 
element, is seen to be of the greatest importance. The 
nitrogen being a harmless, inactive substance, merely 
serves to dilute the oxygen, which would otherwise be too 
vigorous in its action. 

119. Composition of the Atmosphere. The principal 
constituents of the atmosphere are nitrogen, oxygen, 
argon and its related gases, water and carbon dioxid. In 
addition to these there are found ammonia, traces of ozone, 
hydrogen dioxid, nitric acid, hydrogen sulfid, sulfur dioxid, 
and' organic matter. Nearly the whole of the atmosphere 
is made up of the first three constituents. The propor- 
tion of each is practically constant and independent of 
surroundings. Some of the other constituents yary con- 
siderably with the temperature and surrounding condi- 
tions, but their total amount rarely exceeds 0.25 per cent. 
of the whole atmosphere. 

The relative amounts of each of these constituents have 
been very accurately determined, and from samples col- 

LofC. 



100 LECTURE NOTES 

lected in the most widely separated localities ; and while 
very slight variations have sometimes been observed, they 
were the effect of purely local causes. The percentage 
composition of the three principal constituents, both by 
weight and by volume, exclusive of the others, is found to 
be as follows : 

By Weight By Volume 

Oxygen 23.17 20.94 

Nitrogen 75.54 78.12 

Argon, etc., 1.29 .94 



100.00 100.00 

The amount of water, or aqueous vapor, varies greatly 
and depends almost entirely upon the temperature. If a 
cubic meter of air is completely saturated with moisture, 
it is found that at 0° it can hold 4.8 grams of water, while 
at 20° the amount is 17.1 grams, and at 35°, which is 
about our highest summer temperature, the amount is 
39.2 grams. The quantity actually found in the atmos- 
phere is not often more than 60 per cent, of this amount, 
and in very dry climates may not exceed 10 per cent. 
When the atmosphere is cooled, this moisture condenses, 
and falls as rain, snow, or hail. 

The amount of carbon dioxid in the atmosphere varies 
with the locality, the smallest amount being found on the 
sea, and the largest in the neighborhood of large cities 
and town^s. The average amount is about three volumes 
in 10,000, or about 0.03 per cent. 

Ammonia is an important constituent of the atmosphere, 
although the amount found is very small, averaging about 
one part in 1,000,000. The other constituents are due to 
local causes and the amounts found are exceedingly small. 



LECTURE NOTES 101 

120. Argon and its Related Gases. As long ago as 
1785, Cavendish passed electric sparks through a mixture 
of oxygen and air over potassium hydroxid, until no 
further decrease in volume took place. After removing 
the excess of oxygen, he found that a small bubble of gas, 
amounting to about y^- of the volume of nitrogen em- 
ployed, remained unabsorbed. More than a century 
passed before this apparently trifling amount of unab- 
sorbed gas was explained. 

In 1893, Lord Rayleigh, while determining the density 
of nitrogen, discovered that ; ' atmospheric nitrogen," 
that is, nitrogen obtained from the air, was somewhat 
heavier than "chemical nitrogen," or nitrogen which he 
obtained from ammonia. In 1894, Professor Ramsey as- 
sociated himself with Rayleigh in the investigation of this 
difference and soon after they discovered that it was due 
to the presence in the air, of a hitherto unknown gas, 
which they called argon. This discovery was first made 
known in 1895. 

Independent experiments regarding the amount of argon 
in the air gave identical results, and showed that atmos- 
pheric nitrogen, contained 1.186 per cent, of its volume, 
of argon. The density of argon was found to be 19.96, 
so that its molecular weight is 39.9. This is also found 
to be its atomic weight, since examination of its various 
physical properties, shows that its molecule contains only 
one atom. 

Argon is an absolutely inert gas. Every known method 
has been employed to make it combine with some other 
element, but without success, and it is doubtful if any 
compound of argon will ever be known. 

In 1898, Ramsey, while experimenting with liquid argon. 



102 LECTURE NOTES 

discovered the existence of three new gases in the atmos- 
phere. To these he gave the names neon, krypton, and 
xenon. In the summer of 1900, he evaporated a large 
quantity of liquid air, and by fractional distillation of the 
last portions, he succeeded in obtaining these gases in a 
fair state of purity. 

They are all inert gases, and their molecules are all 
monatomic. Their atomic weights, as nearly as can be 
determined, are as follows : neon, 19.9; krypton, 81.7; 
xenon, 128. They exist in small quantities in the atmos- 
phere, but the exact proportions are not yet known. 

The element helium, while not a constituent of the at- 
mosphere, may be mentioned here, since it is closely related 
to these elements, being inert and its molecule monatomic. 

121. The Liquefaction of Gases. — Liquid Air. The 

difference between the liquid and the gaseous state, is 
only a difference in temperature, which may be more or 
less modified by pressure. 

We have already noted that different gases become 
liquids at a low temperature, or under high pressure, or 
by a combination of low temperature and high pressure. 
This is true of all gases, so that air can be liquefied. The 
first gas to be liquefied was chlorin. This was accom- 
plished in 1806, by a pressure of six atmospheres at the 
ordinary temperature. The last of the so-called permanent 
gases were liquefied in 1877, by Cailletet. 

There are two important points which have to be con- 
sidered in the liquefaction of gases. These are the critical 
temperature and the critical pressure. It is found that 
above the point called the critical temperature, a gas 
cannot be liquefied, no matter how great is the pressure; 



LECTURE NOTES 103 

whereas, at a sufficiently low temperature, a gas will 
liquefy under the ordinary atmospheric pressure. The 
critical pressure is the pressure necessary to liquefy a gas 
at the critical temperature. When a gas is at the critical 
temperature and under the critical pressure, the least 
lowering of the temperature or increase of the pressure 
will cause it to liquefy. 

The critical temperature of a gas may be reached by 
cooling it by the evaporation of the liquid formed from 
some easily condensed gas ; or by allowing the gas itself 
to expand suddenly after being very strongly compressed. 
It is the latter principle which is now employed in the 
manufacture of liquid air, which has recently become a 
commercial product, being made on a large scale and at a 
comparatively low cost. 

Liquid air is a light blue liquid. Its critical tempera- 
ture is — 140° and its critical pressure 39 atmospheres. 
It has no constant boiling point since it is not a simple 
substance but a mixture. The oxygen boils at — 181° and 
the nitrogen at — 194°, under ordinary atmospheric pres- 
sure ; so that if liquid air is allowed to stand, the nitrogen 
boils off first and at the end leaves nearly pure oxygen. 

The scientific value of liquid air is very great, because 
of its great refrigerating power, and its relatively low 
cost. What its commercial value may be, is not yet 
determined. 

122. Proof that the Atmosphere is a Mixture. The 

very nearly constant composition of the atmosphere led 
chemists for a long time to suppose that it was a definite, 
chemical compound, of four parts of nitrogen with one of 
oxygen. Investigation, however, shows that this cannot 
be true for the following reasons : 



104 LECTURE NOTES 

1st. All chemical compounds contain their constituent 
elements in exact proportion to their atomic weights. 
The proportion of nitrogen to oxygen is very nearly as 
3.85 to 1, and not as 4 to 1. (Nitrogen and argon are here 
reckoned together.) 

2d. A mixture of the atmospheric elements, in the right 
proportion, has all the properties of the atmosphere, and 
in the mixing there is no variation in temperature, as is 
always the case in chemical combination. 

3d. The proportion of nitrogen to oxygen, while very 
nearly constant, is not absolutely so, as would be the case 
in a chemical compound. 

4th. When air is dissolved in water, the proportion of 
nitrogen to oxygen is quite different from that in the 
ordinary atmosphere, and is exactly in accordance with 
the solubility of the two gases in water. This percentage 
is found to be very nearly, nitrogen, 65, oxygen, 35, in 
the air dissolved in water, which variation would not be 
possible if the atmosphere were a chemical compound. 

123. Ventilation. The processes of combustion and 
respiration require large quantities of oxygen, and at the 
same time produce correspondingly large quantities of 
carbon dioxid. If these processes go on in a closed room, 
the air becomes so charged with the carbon dioxid and 
with the organic matter, which is thrown off by the pro- 
cess of respiration and from the surface of the body, that 
it soon becomes too impure to be breathed with safety. 
It is, as a rule, unsafe to breathe air which contains more 
than 0.10 per cent, of carbon dioxid, and the presence of 
the organic matter materially increases the danger. 

For this reason the subject of ventilation is one of 



LECTURE NOTES 105 

considerable importance to us all, ventilation being the 
process by which foul air is removed and fresh air sup- 
plied. Fortunately for us, living as most of us do in 
poorly ventilated houses, the natural means of ventilation 
are numerous. The old fashioned fireplace is the best of 
ventilators. Stoves, being connected with the chimney, 
also carry off much of the foul air, the necessary fresh air 
coming in through the cracks and crevices about doors 
and windows, as well as through the walls. With our more 
modern heating systems, proper ventilation can best be 
secured by connecting the rooms with a proper ventilating 
shaft, in which a strong draft is maintained. In other 
cases the open window seems to be the only resource. 

It is unfortunate that most of our houses are built with 
no provision for ventilation ; but with proper care we may 
keep our living apartments so well supplied with fresh air, 
that the air we have to breathe will never become 
dangerous. 

COMPOUNDS OF NITROGEN WITH HYDROGEN 

Nitrogen forms three compounds with hydrogen. These 
are : 

Ammonia, NH 3 , 
Hydrazin, N 2 H 4 , and 
Hydrazoic Acid, HN 3 . 

Only the first compound is of much importance. 

Ammonia. Molecular Formula, NH 3 . Density, 8.5. 

124. History and Occurrence. Compounds of ammonia 
were known to the early alchemists, and are often men- 
tioned in their writings, under different names. The 
name ammonia is derived from sal-ammoniac, the name 



106 LECTURE NOTES 

given to common salt found in Egypt, near the temple of 
Jupiter Ammon. This name, sal-ammoniac, was afterward 
applied exclusively to ammonium chlorid. 

Up to the end of the 18th century, ammonia was only 
known in an aqueous solution, which was commonly called 
spirits of hartshorn. In 1774, Priestley first prepared 
the pure gas, which he called alkaline air. 

It occurs in nature in small quantities, but quite widely 
diffused, its compounds being found in the air, in soils, 
and in mineral springs. It is formed by the decay of ni- 
trogenous animal matter, and by electric discharges in 
moist air. The latter gives the compound, ammonium 
nitrate, the action being as follows : 

N 2 +0 + 2H 2 = NH 4 N0 3 . 

Many of the ammonium compounds found in nature, are 
probably derived from this compound. 

125. Preparation. Nitrogen and hydrogen do not unite 
easily to produce ammonia, but can be made to do so by 
passing a silent electric discharge through a mixture of 
these gases. It is most conveniently prepared by heating 
some ammonium salt with calcium oxid (lime), or with 
some other strong base. The following equation repre- 
sents the action when ammonium chlorid and lime are 
used : 

2 NH 4 C1 + CaO == CaCl 2 + H 2 -f 2 NH 3 . 

Ammonia is also produced by the dry distillation of 
animal refuse containing nitrogen, such as bones, hoofs 
and horns. 

Ammonia is produced in large quantities as a by-product 
in the manufacture of illuminating gas. The coal, which 



LECTURE NOTES 107 

is used for this purpose, contains about two per cent, of 
nitrogen, which, in the process of distillation, passes off 
in the form of ammonia ; and this, with other substances, 
is dissolved in the condensed moisture, forming the so- 
called ammojiiacal liquor of the gas works. From this 
liquor, ammonia is obtained in the form of an aqueous 
solution, and so called aqua ammonia. 

126. Properties. Ammonia is a colorless gas. It has 
a characteristic pungent odor, and a sharp, caustic taste. 
It has a very strongly alkaline reaction upon vegetable 
colors. It neutralizes acids, forming with them a series 
of stable compounds called the am?nonium salts. It is one 
of the lightest gases known, its specific gravity being 
only 0.587. 

Ammonia is very soluble in water, the most so of any 
gas known. Under the normal pressure, one volume of 
water at 0° will dissolve 1148 volumes of ammonia, and 
at the ordinary temperature about 750 volumes. If this 
solution be heated to the boiling point, all the gas can be 
driven off. There is probably a weak chemical union be- 
tween the ammonia and the water, forming a compound 
called ammonium hydroxid, and having the molecular 
formula, NH 4 OH. 

Ammonia, under ordinary circumstances, is neither 
combustible nor a supporter of combustion, but if it is 
mixed with oxygen, it will burn with a pale yellow flame. 

Under a pressure of seven atmospheres at the ordinary 
temperature, or at a temperature of — 40° at the ordinary 
pressure, ammonia becomes a colorless liquid. When this 
liquid is allowed to evaporate, a great amount of heat is 
rendered latent. This fact, together with the fact that 



108 LECTURE NOTES 

water dissolves such a quantity of ammonia, has been 
utilized in the construction of Carre's machine for making 
artificial ice. 

127. The Compound Radicals of Nitrogen and Hydro- 
gen. When ammonia combines with an acid to form a 
salt, it does so by addition and not by substitution, thus: 

NH, 4- HC1 = NH CI. 

The union of the ammonia with the atom of hydrogen 
in the acid forms the compound radical, NH^, which is 
called ammonium. This radical acts very much like a 
monad metal, and may be regarded as such. It replaces 
hydrogen in an acid and forms salts, which are called 
ammonium salts. 

There are two other radicals which are common among 
organic compounds, and which are found occasionally in 
inorganic compounds. These are known as the amido 
group, NH 2 , and the imido group, NH. The former is 
univalent, the latter bivalent. 

128. The Halogen Compounds of Nitrogen. Nitrogen 
forms at least one compound with each of the halogens, 
except fluorin. They are among the most violently ex- 
plosive compounds known, and should never be prepared 
except with the greatest care, and in very small quantities. 

Nitrogen chlorid, NCI , is formed when chlorinin excess 
is brought into contact with a warm solution of ammoni- 
um chlorid, thus : 

NH^Cl + 6 CI = NC1 3 + 4 HC1. 

It is a thin, yellowish, oily liquid, with a peculiar, pun- 
gent odor. It evaporates quickly in the air, and the 
vapor attacks the eyes. When it is heated, or when it 



LECTURE NOTES 109 

is brought into contact with phosphorus, or turpentine, 
or even exposed to sunlight, it explodes with very great 
violence. 

Nitrogen bromid, NBr 3 , is a red liquid, and has prop- 
erties similar to the chlorid. 

Nitrogen iodid, NI 3 , or NHI 2 , is formed when iodine is 
brought into contact with a concentrated solution of 
ammonia. It is a black powder, which explodes with 
violence when heated, or touched, or when thrown into 
boiling water. 

COMPOUNDS OF NITROGEN WITH OXYGEN 

There are five compounds of nitrogen with oxygen. 
These are : 

Nitrogen monoxid, or nitrous oxid, N 2 0, 

Nitrogen dioxid, or nitric oxid, NO, 

Nitrogen trioxid, N 2 3 , 

Nitrogen tetroxid, or peroxid, N0 2 , or N 2 4 , and 

Nitrogen pentoxid, N 2 5 . 

Nitrogen Monoxid, or Nitrous Oxid. Molecular For- 
mula, N 2 0. Density, 22. 

129. History and Preparation. Nitrous oxid was first 
prepared by Priestley in 1772. 

It is formed when zinc, and certain other metals, are 
dissolved in very dilute nitric acid. It is most convenient- 
ly prepared by heating the compound ammonium nitrate, 
NH 4 NO g , the action being as follows : 

NH 4 N0 3 = 2 H 2 + N 2 0. 

In preparing nitrous oxid in this way, the ammonium 
nitrate must be heated gently, otherwise the decomposi- 



110 LECTURE NOTES 

tion may take place so rapidly as to cause an explosion. 
Nitrous oxid may also be prepared in a somewhat similar 
way, by heating sodium nitrate with a slight excess of 
ammonium sulfate, thus : 

2 NaN0 3 + (NH 4 ) 2 S0 4 = Na 2 S0 4 + 4 H,0 + 2 N 2 Q. 

130. Properties. Nitrous oxid is a colorless gas having 
a sweetish taste and smell. It is somewhat soluble in cold 
water, a little more than one volume being dissolved at 
0°. It supports combustion almost as vigorously as oxy- 
gen, forming oxids, and liberating nitrogen, the nitrous 
oxid being entirely decomposed at a red heat. 

When nitrous oxid is breathed for a short time, it 
produces a curious kind of nervous excitement or intoxica- 
tion, generally without entire loss of consciousness. This 
effect soon passes away without being followed by any bad 
results. It is for this reason largely employed as an 
ansesthetic in slight surgical operations, especially in 
dentistry, under the name of laughing gas. The anaes- 
thetic effects of this gas were observed by Davy, and this 
substance was the first to be employed for such a purpose 
in a surgical operation. In using it, care should be taken 
to have it perfectly pure. 

Nitrogen Dioxid, or Nitric Oxid. Molecular Formula, 
NO. Density, 15. 

131. Preparation and Properties. This oxid is formed 
when nitric acid acts upon copper, or other metals of a 
like nature, or upon easily oxidizable substances. In this 
action hydrogen is liberated, which, in the nascent state, 
decomposes some of the nitric acid, forming the nitric 
oxid, thus : 



LECTURE NOTES 111 

3 Cu + 8 HN0 3 = 3 Cu(N0 3 ) 2 + 4 H 2 + 2 NO. 

Nitric oxid is a colorless gas, which is only very slight- 
ly soluble in water. It dissolves easily in an aqueous 
solution of ferrous sulfate, imparting a dark brown color 
to the solution. This color is due to the formation of a 
peculiar compound, which has the formula FeSO^NO. If 
this compound is heated, the nitric oxid is entirely ex- 
pelled. This is the principle involved in the well known 
analytical test for nitric acid. 

Nitric oxid is not combustible, but supports the com- 
bustion of certain substances, since, if previously ignited, 
they will continue to burn iu this gas. 

The most characteristic property of nitric oxid is its 
direct union with oxygen, forming nitrogen peroxid, or, 
if the nitric oxid is in excess, nitrogen trioxid. Both of 
these substances have a dark reddish brown color; and so 
if nitric oxid is exposed to the air, it at once becomes 
reddish brown, owing to the formation of nitrogen peroxid. 
It also unites directly with chlorin, forming nitro- 
syl chlorid, NOC1. 

Nitrogen Trioxid. Molecular Formula, N 2 3 . Density, 38. 

132. Preparation and Properties. When easily oxidiz- 
able substances, such as starch, sugar, or arsenious oxid, 
are heated with nitric acid, nitrogen trioxid is formed, 
thus: 

As 2 3 + 2 HNO s = As 2 5 + H s O + N 2 O s . 

It is also formed when four volumes of nitric oxid, and 
one of oxygen, are mixed together. If these vapors be 
led through a freezing mixture, a dark blue liquid is 
formed, which is nearly pure N 2 3 . This substance be- 



112 LECTURE NOTES 

gins to decompose at 0°, forming nitric oxid, and nitro- 
gen peroxid. With water, this compound forms nitrous 
acid, which is very unstable, and only exists at low tem- 
peratures. 

Nitrogen Tetroxid, or Peroxid. Molecular Formula, N0 2 , 
or N 2 4 . Density, above 140 , 23. 

133. Preparation and Properties. This compound may 
be formed by mixing together two volumes of nitric oxid, 
and one volume of oxygen. It is best prepared by heat- 
ing dry lead nitrate, the following equation representing 
the action : 

Pb(N0 3 ) 2 = PbO + O -f- 2 N0 2 . 

Nitrogen peroxid furnishes an interesting example of 
dissociation at ordinary temperatures. If the substance 
is passed through a freezing mixture, it condenses to a 
colorless liquid, which becomes a white solid at — 9°. At 
this temperature its molecular formula is N 2 4 . At 
about 0° it begins to decompose, and the liquid becomes 
yellow. At 22° it boils, and becomes a dark reddish 
brown gas, which grows darker as the temperature is in- 
creased, until it is almost black. At 140° the dissoci- 
ation is complete, the density being 23, which corre- 
sponds to the molecular formula N0 2 . Nitrogen peroxid 
is soluble in water, forming nitric acid and nitric oxid. 

Nitrogen Pentoxid. Molecular Formula, N 2 5 . Molecu- 
lar Weight, 108. 

124. Preparation and Properties. Nitrogen pentoxid 
may be prepared by the action of phosphorus pentoxid on 
nitric acid, thus : 

P,0 6 + 2 HNO, = 2 HPO3 + N 2 6 . 



LECTURE NOTES 113 

It is a colorless, crystalline solid, which melts at 30°, 
and at 47° is entirely decomposed into nitrogen per- 
oxid and oxygen. It is, for this reason, a very pow- 
erful oxidizing agent. It combines with water with great 
energy, giving off much heat, and forming nitric acid. 

135. The Oxy-acids of Nitrogen. Nitrogen forms three 
oxy-acids. These are : 

Hyponitrous acid, H 2 N 2 2 , 
Nitrous acid, HN0 2 , and 
Nitric acid, HN0 3 . 

Only the nitric acid is important. The other two are 
unstable and exist only at low temperatures. They form 
salts which are called hyponitrites and nitrites, respec- 
tively. 

Some of the nitrites are of considerable importance and 
are used quite extensively. They can be made by reduc- 
tion of certain nitrates, and by the oxidation of ammonia. 

Nitric Acid. Molecular Formula, HN0 3 . Molecular 
Weight, 63. 

136. History and Occurrence. Nitric acid has been 
known since the 8th century, having been first made by 
the Arabian alchemist G-eber. It was known to the 
alchemists under the name of aquafortis.. Its true com- 
position was determined by Cavendish in 1785. 

It is produced in small quantities by electric discharges 
in the atmosphere. Its salts, the nitrates, are found in 
quite large quantities, the principal ones being potassium 
nitrate, KN0 3 . called also saltjieter, or niter, and sodium 
nitrate, NaNO g . called Chili saltpeter The former is 
found in large deposits in India, and the latter in parts 



114 LECTURE NOTES 

of Chili and Peru. These deposits were probably formed 
by the oxidation of nitrogenous organic matter, in the 
presence of potassium and sodium bases. 

137. Preparation. Nitric acid is always prepared by 
the action of sulfuric acid on a nitrate, usually of sodium 
or potassium. About equal weights of the nitrate and 
sulfuric acid are mixed in a retort and gently heated, the 
action being represented by the following equation : 

KNCX + H,S(X == KHSO. 4- HNO,. 

6 I I 4 4 ' 3 

If the heat is sufficiently increased, the acid potassium 
sulfate formed, acts upon more of the nitrate, forming 
nitric acid, thus : 

KHS0 4 + KN0 3 = K 2 S0 4 +HNO r 

The nitric acid thus obtained contains water, and has a 
yellow color, from having dissolved some nitrogen per- 
oxid. The water is removed by distilling with concen- 
trated sulfuric acid, and the nitrogen peroxid by gently 
warming the acid, and then leading dry air through it 
until it is cold. This leaves the pure acid. 

138. Properties. Nitric acid, when pure, is a colorless 
fuming liquid. The specific gravity of the concentrated 
acid is 1.53. It is a strong acid, acting on most of the 
metals with great vigor. 

The concentrated acid is an intensely corrosive liquid, 
destroying the skin and causing painful sores. It con- 
verts many organic compounds, such as cotton and gly- 
cerol, into violently explosive substances. The dilute 
acid acts less vigorously but colors the skin, leather, silk, 
wool, and other organic substances, a bright yellow. 

When nitric acid is heated, it begins to boil at 86°, and 



LECTURE NOTES 115 

at the same time is partially decomposed into water, 
nitrogen peroxid, and oxygen. The nitrogen peroxid, 
thus formed, dissolves in it and gives the yellow 
color so common to this acid. By further heating, the 
boiling point continues to rise slowly until it reaches the 
temperature 121°, when it distils without further change 
of temperature, or composition, the strength of acid being 
68 per cent., and the specific gravity 1.42. 

Nitric acid is a very powerful oxidizing agent, owing 
to the ease with which it parts with oxygen. Sulfur and 
phosphorus are oxidized to sulfuric and phosphoric acids, 
respectively, and compounds which exist in a low state of 
valence, are oxidized to a higher state by this acid. Its 
salts, the nitrates, are decomposed by heat, giving 
oxygen, and so are sometimes used as oxidizing agents. 

139. Aqua Regia. Certain of the metals, particularly 
gold and platinum, and also a few compounds, are not 
dissolved by any single common acid. In order to effect 
a solution of these metals, a mixture of three parts of 
hydrochloric and one of nitric acid, is employed. These 
two acids act upon, and decompose each other, with the 
liberation of chlorin, and a volatile compound called nitro- 
syl chlorid, the action being as follows : 

3 HC1 + HN0 3 = 2 H 2 + NOC1 -j-Cl 2 . 

This mixture of acids is called nitro-hydrochloric acid, 
or aqua regia, and the solvent action on the gold or plat- 
inum is due to the fact that the free chlorin formed, 
which is in the nascent state, combines with the metal 
forming a soluble chlorid. It therefore follows, that 
when a metal is dissolved in aqua regia, a chlorid is always 
formed. 



116 LECTURE NOTES 

Phosphorus. Symbol, P. Atomic Weight, 31. 

140. History and Occurrence. Several chemists claim 
the honor of having discovered this element. The best 
evidence, however, seems to prove that it was first pre- 
pared by the alchemist Brand, of Hamburg, in 1669, who 
obtained it from urine, by a process which was kept secret. 
Just a century later, the existence of calcium phosphate 
in bones, was discovered, and in 1771, Scheele discovered 
the method by which it could be prepared from bone-ash. 
This latter process, with slight modifications, is the one 
chiefly employed to-day. 

Phosphorus never occurs free in nature, although in 
combination with other elements it is very widely distrib- 
uted, being found, in quantities more or less minute, in 
almost every substance on the earth's crust. 

The most important compound of phosphorus is calcium 
phosphate, Ca 3 (P0 4 ) 2 . This is found in immense quanti- 
ties in the so-called phosphate beds of South Carolina, in 
the guano beds of the Caribbean islands, and in a number 
of minerals. It occurs in small quantities in the rocks, 
and by the decomposition of these finds its way into the 
soil. From the soil it is taken up by the plants, where 
it forms an essential constituent of the fruit and seeds. 

From the plants phosphorus finds its way into the ani- 
mal body, where it is found in the brain and tissues, and 
especially in the bones, which consist of about 50 per 
cent, of calcium phosphate. When the tissues waste 
away, they are replaced by fresh material, the phos- 
phorus being excreted in the urine, chiefly in the form 
of sodium ammonium phosphate, or microcosmic salt, 
Na(NH 4 )HP0 4 , 4 H a O. 



LECTURE NOTES 117 

141. Preparation. Phosphorus is nearly always pre- 
pared from bones. The bones, which contain a large 
amount of organic material, are first heated in a retort. 
The organic material is decomposed, giving off ammonia 
and some other gases, together with a liquid called bone- 
oil. The residue which is known as bone-black or animal- 
charcoal consists chiefly of calcium phosphate and carbon. 
The bone black is first used by the sugar refiners for clar- 
ifying the syrup, and after it has become useless for this 
purpose, it is burned in an open fire, forming bone-ash, 
which is mostly calcium phosphate. 

The bone-ash is treated with sufficient sulfuric acid 
(specific gravity, 1.60) to decompose it, and to convert all 
the calcium into insoluble calcium sulfate, leaving phos- 
phoric acid in solution, thus : 

Ca 3 (P0 4 ) 2 + 3 H 2 SO + = 3 CaSO, + 2 H s P0 4 . 

The insoluble calcium sulfate is separated from the rest 
by filtering, and the solution, which contains the phos- 
phoric acid, evaporated to a thick syrup. The acid is 
then mixed with charcoal or coke and thoroughly dried 
in cast-iron pots. This changes the common or orthophos- 
phoric acid into metaphosphoric acid, thus : 

H 8 PO, = HP0 3 + H 2 0. 

The dry mass, consisting of metaphosphoric acid and 
carbon, is then heated to bright redness in earthenware 
retorts, when the acid is decomposed, thus : 

4 HP0 3 -f 12 C = 2 H 2 + 12 CO + P 4 . 

The mixture of gases and phosphorus vapor passes 
through an earthenware pipe into a suitable vessel and 
under water, where the phosphorus vapor is condensed. 



118 LECTURE NOTES 

The crude phosphorus thus obtained is purified by 
mixing with potassium bichromate and sulfuric acid, 
which oxidizes the impurities, or by redistillation and 
pressing through wash leather, after which it is cast into 
sticks, in copper tubes, all the operations being carried 
on under water. 

142. Properties. Ordinary phosphorus is a slightly 
yellow, or almost colorless, crystalline solid. If it solid- 
ifies slowly it is almost perfectly transparent, but if it 
cools quickly it is only translucent. If it is kept for some 
time, even in the dark, it becomes coated with an opaque 
white film, and if it is exposed to the light it soon becomes 
yellow, then red-brown, and finally becomes almost black. 

At a low temperature phosphorus is brittle, but at 
about 15° it becomes soft and wax-like, and is easily cat 
with a knife, although it still shows a crystalline struc- 
ture. It has a specific gravity of 1.83. 

Phosphorus melts at 44°, and, if heated in an atmos- 
phere free from oxygen to a temperature of 290°, it boils, 
yielding a colorless vapor, which has a specific gravity of 
4.3. This corresponds to a density of 62 and a molecular 
weight of 124, which shows that its molecule consists of 
four atoms. 

Phosphorus is very nearly insoluble in water, but dis- 
solves very easily in carbon disulfid. It is also somewhat 
soluble in alcohol, chloroform, olive oil, and many organic 
solvents. If a solution in carbon disulfid is allowed to 
evaporate, the phosphorus is deposited in the form of 
colorless, isometric crystals. 

If a piece of phosphorus is placed in a vacuous tube and 
kept in the dark, the phosphorus will slowly volatilize, 



LECTURE NOTES 119 

and will be deposited on the sides of the tube in the form 
of bright colorless crystals. 

Phosphorus is an exceedingly inflammable substance. 
It takes fire at a temperature a little above its melting 
point, and burns with a bright flame, forming phosphorus 
pentoxid P 2 5 . If a piece of dry phosphorus is exposed 
to the air, it at once begins to oxidize, and the heat of 
this oxidation is sufficient to cause it to take fire after a 
short time. It is, for this reason, always kept under 
water. 

Phosphorus is a violent poison, and when taken into 
the system through the stomach, probably acts upon the 
blood, depriving it of oxygen. There is another form of 
slow poisoning which is produced by continued inhaling 
of phosphorus vapors or the white fumes from the slow 
oxidation of phosphorus. This is mostly confined to 
persons engaged in the manufacture of phosphorus and 
articles containing phosphorus, such as matches, and 
general^ manifests itself in the form of necrosis of the 
bones of the jaw and nose, and of caries of the teeth. No 
good antidote for phosphorus poisoning is known, perhaps 
the best is the breathing of oxygen. 

When phosphorus is exposed to moist air, it oxidizes 
slowly, and, in the dark, appears luminous, emitting a 
greenish-white light. At the same time it gives off a 
white vapor, which has a garlic-like odor, and is poison- 
ous. This phenomenon, which is known as phosphores- 
cence., is one of the most characteristic properties of 
phosphorus. If a few drops of a solution of phosphorus 
in carbon disulfid are poured on to a piece of filter paper, 
the carbon disulfid soon evaporates and the paper shows 
bright phosphorescence in the dark. The thin film of 



120 LECTURE NOTES 

phosphorus on the filter paper oxidizes so rapidly that it 
soon takes fire and burns. 

If phosphorus is placed in boiling water a small portion 
of it volatilizes with the steam. If the steam passes 
through a cod denser, a phosphorescent ring will be seen 
at the point where the steam begins to condense. ; This 
phenomenon is utilized for the detection of phosphorus in 
cases of poisoning. 

143. The Allotropic Forms of Phosphorus. There are 
three allotropic forms of phosphorus, which differ from 
each other in a very marked degree. 

The first form is the common or octohedral phosphorus, 
and is a colorless or slightly yellow crystalline solid. Its 
properties have already been given. 

The second form is produced when the common phos- 
phorus is heated to about 250° in a gas which has no 
action upon it, or in a closed vessel to 300°. By this 
action it is changed into a red-brown amorphous solid, 
which is known as red phosphorus, or amorphous phos- 
phorus, and which has a specific gravity of 2.11. Red 
phosphorus, when pure and dry, has neither odor nor 
taste, but in moist air it is very slowly oxidized. It is not 
poisonous, is not luminous in the dark, and does not take 
fire until heated to about 260°. When heated in the 
absence of air to 350°, it slowly changes back to the ordi- 
nary variety. 

The third form is produced when red phosphorus is 
heated under pressure to 580°, or, when phosphorus, 
in contact with metallic lead, is heated for some hours in 
sealed tubes at a very high temperature. This form is a 
dark red crystalline mass which is known as metallic or 



LECTURE NOTES 121 

rhombohedral phosphorus. It has a specific gravity of 
2.34, and conducts electricity. 

Other so-called modifications have been described, but 
their existence is doubtful. 

144. Uses. Phosphorus is used quite extensively in the 
laboratory, where it is employed in making certain com- 
pounds which are much used in the preparation of the 
anilin dyes. 

The principal use of phosphorus in the arts is in the 
manufacture of matches. The bits of wood are first tipped 
with paraffin or, as is sometimes the case, with sulfur. 
and then with a paste made of phosphorus, with either 
potassium chlorate or nitrate, and glue. These ignite 
when rubbed upon any rough surface. The so-called 
safety matches are tipped with a mixture of potassium 
chlorate, potassium bichromate, red lead, and antimony 
sulfid. Under ordinary circumstances these will not 
ignite until rubbed upon a prepared surface, which con- 
sists of a mixture of red phosphorus and antimony sulfid. 
They can be inflamed, however, by striking them very 
quickly on a smooth, non-conducting surface, such as 
glass, paper, etc. 

COMPOUNDS OF PHOSPHORUS WITH HYDROGEN 

There are three compounds of phosphorus and hydrogen 
known. They are distinguished by the state of aggrega- 
tion in which they exist, as 

Gaseous hydrogen phosphid, or Phosphin, PH 3 , 
Liquid hydrogen phosphid, P 2 H 4 , and 
Solid hydrogen phosphid, Pfi 2 , 



122 LECTURE NOTES 

Phosphin. Symbol, PH 3 . Density, 17. 

145. Preparation and Properties. This substance is 
best prepared by heating phosphorus in a solution of 
sodium hydroxid, when phosphin and sodium hypophos- 
phite are formed, thus : 

3 NaOH + P 4 4- 3 H 2 == 3 NaH 2 P0 2 + PH S . 

Phosphin is also formed when calcium phosphid is 
thrown into water. 

Phosphin is a colorless gas. It has a very disagreeable 
odor somewhat like rotten fish. The pure gas takes fire 
at a temperature above 100° ; but the gas as prepared by 
either of the above methods, contains traces of the liquid 
hydrogen phosphid, which renders it spontaneously in- 
flammable, so that if it be allowed to pass through water, 
each bubble as it reaches the surface ignites with a slight 
explosion, forming a beautiful ring of phosphorus pent- 
oxid, which shows a remarkable vortex motion. 

Phosphin is a very poisonous gas, which acts upon the 
blood. If it is present in the air, even in small propor- 
tions, it soon causes difficulty in breathing, and death. 
It is a weak base, and forms compounds which are analo- 
gous to those formed by ammonia, and which are called 
phosphonium compounds. 

The liquid and solid hydrogen phosphids are not of 
especial importance. 

146. The Halogen Compounds of Phosphorus. Phos- 
phorus forms compounds with each of the halogen elements, 
in most cases by direct union, and sometimes the union 
takes place with explosive violence. The chlorids are the 
most important of these compounds. 



LECTURE NOTES 123 

If dry chlorin is led over an excess of phosphorus, phos- 
phorus trichlorid, PC1 8 , is formed. 

The trichlorid is a colorless liquid, which fumes strongly 
in the air, and boils at 76°. When brought into contact 
with water, it decomposes into hydrochloric and phos- 
phorous acid thus : 

PC1 3 -f 3 H 2 = 3 HC1 + H 3 PO s . 

If in making the trichlorid, the chlorin is in excess, or 
if chlorin be led into the trichlorid, phosphorus penta- 
chlorid is formed. 

Phosphorus pentachlorid is a yellowish white crystalline 
solid, which is much used in organic chemistry for replac- 
ing the hydroxyl group with chlorin. 

Phosphorus forms two bromids, two iodids, two 
fluorids, and a number of mixed, or double compounds. 

COMPOUNDS OF PHOSPHORUS WITH OXYGEN 

147. The Oxids of Phosphorus. There are four oxids 
of phosphorus known. These are : 

Phosphorus suboxid, P 4 0, 
Phosphorous oxid, P 4 6 , 
Phosphorus tetroxid, P 2 4 , and 
Phosphorus pentoxid, P 2 5 . 

The suboxid and tetroxid are of little value, the former 
being a red amorphous powder, while the latter forms 
colorless rhombic crystals. 

148. Phosphorous Oxid, or Phosphorus Trioxid, P 4 6 * 

Phosphorous oxid is formed when phosphorus is allowed 
to oxidize in a limited supply of air. Care must be taken 



124 LECTURE NOTES 

that the temperature does not get too high, or the pent- 
oxid will be formed. 

It is a white crystalline solid, with an odor like garlic, 
and, when dry, does not redden litmus. It dissolves 
slowly in cold water forming phosphorus acid. With hot 
water the action is quite violent, phosphin, phosphoric 
acid and red phosphorus being formed. If exposed to the 
air, it slowly oxidizes to the pentoxid. It was at first 
supposed that the molecule was P 2 3 , but determination 
of its vapor density shows that the molecule is P 4 6 - If 
heated in a sealed tube to 440° it forms the tetroxid and 
red phosphorus. 

149. Phosphorus Pentoxid, P 2 5 . Phosphorus pent- 
oxid is always formed when phosphorus burns in the air, 
or in oxygen. To obtain this compound, a piece of phos- 
phorus may be burned on a plate under a bell jar. 

Phosphorus pentoxid is a white amorphous powder. 
When pure it has neither color nor odor, but it may con- 
tain traces of the trioxid, which gives it a garlic like odor, 
or some red phosphorus, or some of the sub-oxid, which 
gives it a reddish color. It is very deliquescent, and, if 
exposed to the air, soon becomes liquid, forming metaphos- 
phoric acid, thus : 

P,0 5 + H.O .= 2 HPO,. 

When perfectly dry it has no action on dry litmus 
paper. It dissolves in water with a hissing sound and 
gives off much heat. 

Phosphorus pentoxid is used in the laboratory for des- 
icating purposes and for removing the elements of water 
from compounds which contain them. [See Nitrogen 
Pentoxid, Page 112.] 



LECTURE NOTES 125 

COMPOUNDS OF PHOSPHORUS WITH OXYGEN AND 
HYDROGEN 

150. The Oxy-acids of Phosphorus. There are six 
oxy-acids of phosphorus, although two of them are deriv- 
atives of one of the others. These are : 

Hypophosphorous acid, H 3 P0 2 , 
Phosphorous acid, H 3 P0 3 , 
Hypophosphoric acid, H 4 P a 6 , 
Phosphoric acid, H 3 P0 4 , 
Pyrophosphoric acid, H 4 P 2 7 , and 
Metaphosphoric acid, HPO s . 

The last two are derivatives of phosphoric acid. 

None of the first three are very important acids, and 
the second and third are quite unstable. Their salts are 
quite stable and of more importance. 

151. Phosphoric Acid, H 3 P0 4 , and its Derivatives. Phos- 
phoric acid, which is also known as orthophosphoric acid, 
may be prepared in many ways, but is best prepared by 
the oxidation of phosphorus by means of nitric acid. Red 
phosphorus is commonly employed in this reaction, which 
is as follows : 

P 3 + 5 HNO s -f 2 H 2 = 3 H 3 P0 4 + 5 NO. 

It may also be prepared by dissolving phosphorus pent- 
oxid in hot water. Commercially it is prepared from 
bone-ash. 

Phosphoric acid is a white crystalline solid, having no 
odor, but an agreeable acid taste. It is very soluble in 
water and forms salts which are called phosphates. It is 
a tribasic acid and so there are three classes of phosphates, 
two acid and one normal. These are also called primary, 



126 LECTURE NOTES 

secondary and tertiary phosphates, according as one, two, 
or three atoms of hydrogen are replaced by a metal. 

If phosphoric acid is heated to 215°, it loses water and 
forms pyrophosphate acid, thus : 

2H 3 P0 4 = H ( P ! 7 + H 2 0. 

Pyrophosphoric acid is a tetrabasic acid. Its salts 
which are called pyrophosphates, are formed by heating 
the secondary phosphates. 

If either phosphoric, or pyrophosphoric acid is heated 
to 400°, more water is given off, and there is formed meta- 
phosphoric acid, thus : 

H 3 P0 4 = HP0 3 + H 2 0. 

There are also a number of polymeric metaphosphites, 
which may be regarded as salts of certain hypothetical, 
polymeric forms of this acid. 

Metaphosphoric acid is a glassy, transparent solid, 
which is often called glacial phosphoric acid. It melts at 
a red heat, but does not decompose. It is a strong, mono- 
basic acid, and its salts may be formed by heating the 
primary phosphates. 

152. The Constitution of the Acids of Phosphorus. 

The constitution of these six acids has been the subject of 
much discussion, and the questions are not yet settled. 
Some would regard them all as derived from phosphin 
PH 3 , in which case they would all contain trivalent phos- 
phorus. The hypophosphorus and phosphorous acids 

it r\ 

would have as their constitutional formulas tt n > P — H, 

H O 
and tt r)]>P — O H, respectively ; but the orthophos- 



LECTURE NOTES 127 

phoric acid would have the constitutional formula 

XT Q 

-p. q>P — O — OH, and the derivative acids would have 

a similar constitution, which would hardly be in accord 
with the relative stability of these acids. 

No salt of hypophosphorous acid is known in which 
more than one atom of hydrogen in the acid is replaced 
by the metal. The acid is, therefore, monobasic. It 
therefore contains but one hydroxyl group, while if the 
phosphorus were trivalent it must contain two. 

A trivalent salt of phosphoric acid is known, viz. : 
Na 3 P0 3 . This salt is very unstable and is decomposed 
by water, forming sodium phosphite, Na 2 HPO s , which is 
generally regarded as the normal sodium phosphite. 

The phosphoric acids undoubtedly contain pentavalent 
phosphorus, and we shall get the clearest idea of the 
constitution of all the acids, if we suppose that each one 
contains the trivalent phosphoryl group =(P=0). All 
the phenomena observed in connection with these acids 
and their salts, can be accounted for by this supposition. 
The constitutional formula of each acid would then be as 
follows : 

HO>_ 
HO\ HO\ HO/ 

(1.) H— P = 0, (2.)HO-P = 0, (3.) | 

H/ H / H0 ^P-0 

HO/- U 

HO\ 
HO\ HO— P = HO\ 

(4.)HO-P=0, (5.) 0< , (6.) H=P=0. 

HO^ HO— P=0 

HO/ 

These numbers correspond to the list on page 125. 



128 LECTURE NOTES 

COMPOUNDS OF PHOSPHORUS WITH SULFUR 

153. The Phosphorus Sulfids. It was for a long time 
believed that there were two series of compounds of these 
two elements, which were called sulfur phosphids, and 
phosphorus sulfids, respectively. The former series 
proved to have been only solutions of sulfur in phosphorus. 

Of the latter series, there are several compounds known. 
They are all formed by direct union of the elements, but 
unless great care is taken, the union is often attended by 
a violent explosion. 

Phosphorus pentasulfid P 2 S., is the only one of much 
importance. This is a gray solid, which melts at about 
275°, and boils at 518°. It is used in organic chemistry 
for forming sulfur compounds, which it does by replacing 
oxygen with sulfur. With alcohol it forms mercaptan, 
thus : 

5 C 2 H 5 OH + P 2 S 5 = 5 C 2 H 5 SH + P 2 O s 

Water decomposes phosphorus pentasulfid, forming 
phosphoric acid, and hydrogen sulfid. 

Arsenic. Symbol, As. Atomic Weight, 75. 

154. History and Occurrence. Certain compounds of 
arsenic, especially the native sulfids, called realgar and 
orpiment, were known to the ancients, and arsenious oxid 
was known to most of the alchemists as early as the 
eighth century. 

The alchemists valued these compounds of arsenic very 
highly. They knew that copper became white when 
heated with these arsenic compounds, and believed this 
change in color, to be due to a partial transmutation of 
copper into gold. Arsenic was first obtained, and recog- 



LECTURE NOTES 129 

nized as an element, in the latter part of the eighteenth 
century. 

Arsenic occurs free in nature, in small quantities, and 
in a few localities. In combination with other elements, 
it is quite widely distributed. It occurs combined with 
sulfur, and oxygen, and in compounds with the metals, 
called arsenids. Small quantities of arsenic occur in con- 
nection with certain metallic sulfids, such as pyrite, and 
this is one of its important sources. 

155. Preparation. If arsenical pyrite, is heated, the 
arsenic is driven off, and ferrous sulfid remains, thus : 

FeSAs = FeS + As. 

If ores containing arsenic, are roasted in a furnace, so 
constructed as to admit a plentiful supply of air, arsen- 
ious oxid is formed. This substance, which is q uite volatile, 
is collected in suitable chambers. The arsenious oxid is 
then mixed with charcoal, and heated in earthenware 
tubes, when the free arsenic volatilizes, and collects in 
the cooler portion of the tubes. The action is as follows : 

As 4 6 -f 6 C = 6 CO + As 4 . 

In order to purify the crude arsenic, a little charcoal is 
added and it is resublimed. 

156. Properties. Arsenic is a steel-gray, crystalline 
solid, having a specific gravity of 5.72. When heated 
under ordinary conditions it does not melt, but passes at 
once from the solid to the gaseous state. Under pressure, 
arsenic fuses at about 500°. The vapor of arsenic has a 
lemon-yellow color, and a strong garlic-like odor, proba- 
bly due to a partial oxidation of the element. When 
heated in the air it takes fire, and burns with a bright 
white flame. The vapor of arsenic has a density of 150, 



130 * LECTURE NOTES 

which shows that the molecule contains four atoms. Ar- 
senic and most of its compounds are very poisonous. 
Arsenic exists in three allotropic modifications. If it is 
made to sublime in a stream of hydrogen gas, the portion 
which is deposited nearest the heated portion of the tube 
will be steel-gray and crystalline. A little further on it 
will be a black glittering mass, and still further on a 
light yellow powder. The last two varieties are amor- 
phous. The black mass has a specific gravity of 4.71 and 
the yellow powder of 3.70. Both these modifications 
change into the first, or crystalline form, when heated to 
360°. 

The relations of arsenic to the other elements are pecu- 
liar. In the free state it is quite metallic in its physical 
properties. In its compounds with the non-metallic 
elements it very closely resembles phosphorus, while with 
the metals it is analogous to sulfur, since the two ele- 
ments can mutually replace each other. 

COMPOUND OF ARSENIC WITH HYDROGEN 

Hydrogen Arsenid, or Arsin. Molecular Formula, AsH 8 . 
Density, 39. 

157. Preparation. Arsin is formed whenever an arsenic 
compound is acted upon by nascent hydrogen. It is most 
conveniently prepared in this way, by introducing a solu- 
tion containing arsenic into a hydrogen generator. The 
action upon arsenious oxid is as follows : 

As 4 O g -f 12 H 2 = 4 AsH 3 + 6 H 2 0. 

If prepared in this way, arsin always contains consid- 
erable hydrogen. In order to prepare the pure gas, an 



LECTURE NOTES 131 

alloy of zinc and arsenic is acted upon by dilute sulfuric 
acid, thus : 

Zn 3 As 2 -f 3 H 2 S0 4 = 3 ZnS0 4 -f 2 AsH 3 . 

Arsin is also formed when arsenic compounds come in 
contact with organic matter which is undergoing decom- 
position. It is in this way that cases of arsenical 
poisoning sometimes occur in dwelling houses. Arsenic 
compounds were formerly, and are still occasionally, 
found in the pigments used in wall-papers. If from 
dampness, or other causes, the walls become mouldy, 
nascent hydrogen is evolved and arsin formed. 

158. Properties. Arsin is a colorless gas, with a garlic- 
like odor. It is inflammable and burns in the air with a 
bluish-white flame, forming arsenious oxid, As 4 O g . 

Arsin is easily decomposed by heat into its elements, 
and so if the gas is made to pass through a glass tube 
which is heated at one point, black amorphous arsenic is 
deposited on the tube just beyond the heated point, as a 
shining mass, which is known as the arsenic mirror. This 
fact is employed in the well known Marsh's test for arsenic. 

Arsin is an extremely poisonous substance, a single 
bubble of the pure gas having been known to produce fatal 
results. The very greatest care should therefore be ob- 
served while experimenting with this compound. 

159. The Halogen Compounds of Arsenic. Arsenic 

forms one compound with each of the halogen elements, in 
which it is trivalent. With iodin it also forms two other 
unstable compounds, a di-iodid and a penta-iodid. The 
compounds combine directly and show quite a strong 
affinity for each other. 



132 LECTURE NOTES 

Arsenic trichlorid, AsCl 3 . may be formed by passing 
dry chlorin over heated arsenic. It is purified by distil- 
ling with excess of arsenic. 

It is most conveniently prepared by heating arsenious 
oxid with sulfuric acid and sodium chlorid. The last two 
compounds form , hydrochloric acid which acts upon the 
arsenious oxid, thus : 

As 4 6 + 12 HC1 = 6 H z O -f 4 AsCl 3 . 

The water formed remains with the sulfuric acid, and 
the trichlorid distils over. 

Arsenic trichlorid is a colorless, oily liquid, which boils 
at 130°. When exposed to the air it gives off dense white 
fumes. It is extremely poisonous. It is decomposed by 
water, forming arsenious oxid and hydrochloric acid. 

The other halogen compounds of arsenic are quite sim- 
ilar in their chemical properties. They are not of espe- 
cial importance. 

COMPOUNDS OF ARSENIC WITH OXYGEN 

Arsenic forms two compounds with oxygen. These are : 

Arsenious oxid, As O , and 
Arsenic pentoxid, As 2 5 . 

Each of these oxids is an acid anhydrid, since it forms 
an acid with water. 

160. Arsenious Oxid, As 4 6 . This substance is known 
commercially as white arsenic or, more commonly, simply 
as arsenic. It is found in nature as the mineral arsenolite. 
It is formed when arsenic is burned in the air or oxygen, 
or when arsenical ores are roasted in a current of air. 



LECTURE NOTES 133 

The arsenious oxid is volatile, and, with other gases, 
passes through what are called poison chambers, where it 
is condensed as a white powder, which is purified by re- 
sublimation. 

Arsenious oxid is a white, crystalline solid. It is odor- 
less but has a sweetish, metallic taste. It is only slightly 
soluble in water, but more soluble in hydrochloric acid. 
With the soluble bases it dissolves and combines chemi- 
cally, forming salts of arsenious acid. When it is allowed 
to crystallize from a solution, or when its vapor is allowed 
to condense on a cold surface, it forms regular octahed- 
rons. It has a specific gravity of 3.74, and, if heated to 
218°, volatilizes without fusing, giving a colorless vapor. 
This vapor has a density of 198, which shows that its 
molecule is As 4 6 . It can be fused under pressure. 

If arsenious oxid is fused and then allowed to solidify, 
at a temperature near the fusing point, it forms a color- 
less, transparent, amorphous mass, which is an allotropic 
modification. This soon becomes white, and opaque, like 
porcelain, and gradually changes back to the crystalline 
form. 

A third form, consisting of rhombic prisms, is obtained 
when a boiling solution of potassium hydroxid is satura- 
ted with arsenious oxid, and the latter allowed to crystal- 
lize. It is also formed when the common variety, heated 
in a closed tube to 400°, is allowed to cool slowly. 

Arsenious oxid is the compound from which most other 
arsenic compounds are made. It is also employed in the 
manufacture of most commercial articles that contain 
arsenic. It is employed to some extent in medicine, and 
is the compound most always found in cases of accidental 



134 LECTURE NOTES 

or intentional poisoning. It is a singular fact, that per- 
sons can accustom themselves to the use of this substance, 
so as to be able to withstand doses which would prove 
fatal, if given to an ordinary person. Its continued use 
is, however, very dangerous. The best antidote for poison- 
ing with arsenic, is freshly precipitated ferric hydroxid. 

161. Arsenic Pentoxid, As 2 5 . When arsenic is burned 
in the air or in oxygen, arsenious oxid is formed, and not 
the pentoxid, as is the case when phosphorus is burned. 

Arsenic pentoxid is obtained by oxidizing arsenious 
acid with nitric acid, and heating to redness, the arsenic 
acid, thus produced, thus : 

2 H 3 As0 4 = 3 H 2 + As 2 5 . 

Arsenic pentoxid is a white, deliquescent solid, which 
dissolves slowly in water, forming arsenic acid. When 
very strongly heated it breaks up into arsenious oxid 
and oxygen. 

COMPOUNDS OF ARSENIC WITH OXYGEN AND 
HYDROGEN 

162. The Oxy acids of Arsenic. There are four oxy- 
acids of arsenic, which are analogous to four of the acids 
Of phosphorus. They are : 

Arsenious acid, H 3 As0 3 , 
Arsenic acid, H 3 As0 4 , 
Pyro-arsenic acid, H 4 As 2 T , and 
Metarsenic acid, HAsO s . 

Arsenious acid, H 3 AsO s , is formed when arsenious oxid 
dissolves in water. Its soluble salts are formed by dis- 
solving arsenious oxid in the corresponding base. It is a 
weak, tribasic acid, existing only in an aqueous solution. 



LECTURE NOTES 135 

Some of its salts, which are called arsenites, are quite 
important bodies. The copper salt, CuHAsO s , is the well 
known pigment called ScheeWs green. 

Arsenic acid, H AsO. called also ortho-arsenic acid, is 
formed when arsenious oxid is oxidized in the presence of 
water, thus : 

As.O, + 4 HNO, + 4 H,0 = 4 H 3 AsO t + 2 N 2 O s . 

Arsenic acid is used in calico printing and in the man- 
ufacture of certain of the anilin dyes. Its salts, the 
arsenates, correspond very closely to the phosphates. 

Pyro-arsenic acid, and metarsenic acid, are each formed 
by heating arsenic acid, the former to 180°, and the latter 
to 200°. They each dissolve in water and are at once 
transformed into the ortho-acid. Their salts act in a 
similar way. 

COMPOUNDS OF ARSENICjWITHj SULFUR 

Arsenic forms three compounds with sulfur. These are : 

Arsenic disulfid, As 2 S 2 , 
Arsenic trisulfid, As 2 S 3 , and 
Arsenic pentasulfid, As 2 S 5 . 

163. Arsenic Disulfid, As 2 S 2 . This compound, which 
was known to the ancients, is found in nature as the min- 
eral Realgar, forming ruby-red crystals. It may be 
formed artificially by fusing together; arsenic and^sulfur 
in the right proportions. This product is known as ruby 
sulfur, and is employed in the arts for the manufacture 
of the so-called Indian-, or ^white-fire, which consists of 
two parts of arsenic disulfid and twenty-four pa**ts of 
potassium nitrate. When ignited, this mixture burns 
with a splendid white light. 



136 LECTURE NOTES 

This compound, as well as many other arsenic com- 
pounds, was formerly much used as a pigment ; but on 
account of its poisonous qualities it is no longer used to 
any extent. 

164. Arsenic Trisulfid, As 2 S 3 . This compound, which 
was also known to the ancients, is found in nature as the 
mineral Orpirne?it, forming bright yellow crystals. It is 
formed when hydrogen sulfid, is passed through an acid 
solution of arsenious oxid, thus : 

As 4 O e + 6 H 2 S = 6 H 2 + 2 As 2 S 3 . 

An impure compound, known commercially as King's 
yellow, and which is a mixture of arsenious oxid and tri- 
sulfid, is formed when arsenious oxid and sulfur are 
heated together. 

Arsenic trisulfid is used in the arts in the printing of 
indigo colors. A mixture of this compound with calcium 
hydroxid and water, under the name of Rusma is used by 
tanners, in the East, for removing hair. 

165. Arsenic Pentasulfid, As 2 S 6 . If hydrogen sulfid is 
led through a hot solution of arsenic acid, there is formed 
a mixture of arsenic trisulfid and sulfur. The pentasulfid 
can be formed by fusing together arsenic trisulfid and 
sulfur, or better, by decomposing some sulfarsenate with 
dilute hydrochloric acid, thus : 

2 K 3 AsS 4 + 6 HC1 = 6 KC1 + 3 H 2 S + As 2 S g . 

It is a bright yellow powder which can be sublimed 
without decomposition. 

166. The Sulfo-Salts, or Thio-Salts. If certain sul- 
fids, such as those of arsenic, antimony, tin, gold, and 
platinum, are treated with a solution of some alkaline 



LECTURE NOTES 137 

sulfid, such as potassium sulfid, K 2 S, or ammonium sulfid, 
(NH 4 ) 2 S, they dissolve, with the formation of a peculiar 
class of compounds, which have the same relation to the 
sulfids, as do the common salts to the oxids. These com- 
pounds are true salts, but have the oxygen replaced by 
sulfur, and in order to distinguish them, they are called 
sulfo-salts, or thio-salts. Thus, if arsenic trisulfid be dis- 
solved in a solution of potassium sulfid, potassium sulfar- 
senite is formed, thus : 

As 2 S 3 + 3 K 2 S = 2 K s AsS 3 . 

In the same way the pentasulfid forms potassium sulf- 
arsenate, K 3 AsS 4 . 

The alkaline sulfo-salts are soluble in water, and are 
decomposed by dilute acids. The sulfo-salts of the other 
metals are insoluble in water. 

167. Marsh's Test for Arsenic. Since arsenic and all 
of its compounds are poisonous, and since these compounds 
are used quite extensively in the commercial world, cases 
of poisoning by arsenic, accidental or intentional, are not 
uncommon. It is, therefore, important that the chemist 
should be familiar with the best methods for detecting 
the presence of arsenic in such cases. A number of 
methods for the detection of arsenic have been devised, 
one of the best, and most commonly used, is known as 
Marsh's test. 

Before the suspected substance is actually tested by 
this, or any other method, it must be freed from all organic 
matter. This is best done by oxidixing the whole with 
hydrochloric acid and potassium chlorate. The whole is 
then filtered, and the solution, which contains the arsenic 
as arsenic acid, evaporated nearly to dryness. 



138 LECTURE NOTES 

A portion of this solution, which may be diluted, if 
necessary, is then placed, in a small hydrogen generator 
containing pure zinc and pure sulfuric acid. If arsenic 
is present, the nascent hydrogen will combine with it 
forming arsin. The gases are then made to pass through 
a hard glass tube heated in one or more places, where 
arsin, if present, is decomposed, and the arsenic mirror 
formed on the inside of the tube. If the gas which issues 
from the end of the tube is ignited, it will burn with a 
bluish-white flame, and if a piece of cold porcelain be held 
against the flame, a black spot of metallic arsenic is in- 
stantly formed on the porcelain. By exercising proper 
care in this operation, exceedingly minute traces of arsenic 
can be detected. 

Compounds of antimony treated in this same way give 
a similar result. It is easy to distinguish one from the 
other, however, for arsenic is quite black and the spot on 
porcelain has a bright metallic luster, while antimony is 
gray-black and the spot on porcelain is dull and has little 
or no luster. Arsenic dissolves very easily in a solution 
of bleaching powder, while antimony is insoluble. These, 
and many other tests, may be employed to confirm the 
presence or absence of arsenic. 

1G8. Reinsch's Test for Arsenic. This is another ex- 
cellent method for detecting arsenic, and in some ways is 
preferable to that of Marsh. The best results are obtained 
from arsenious compounds, so that arsenic compounds 
should be reduced before applying this test. 

To the suspected liquid add about one eighth of its 
volume of hydrochloric acid and a small piece of bright 
copper-foil and boil. If arsenic is present it will be de- 
posited upon the copper as a gray film, which is an alloy 



LECTURE NOTES 139 

of copper and arsenic. If very small quantities of arsenic 
are present, the film may appear bluish or violet, in which 
case the boiling should be continued for some time. If a 
piece of this foil is placed in a small hard glass tube, 
drawn out into a capillary neck, and heated, a narrow 
ring of octahedral crystals of arsenious oxid will be found 
in the capillary neck. These may be examined under the 
microscope, and the presence of arsenic thus confirmed. 

169. The Carbon Group. This group comprises the 
non-metallic elements carbon and silicon, together with 
the metals tin and lead, and some rare elements. The 
non-metals are quadrivalent and very rarely bivalent, 
while the metals are usually bivalent, and occasionally 
quadrivalent. Carbon and silicon are among the most 
important elements known, one or the other being found 
in nearly every naturally occurring compound. Carbon 
is an important constituent of all animal and vegetable 
compounds, while silicon bears a similar relation to most 
of the rocks and soils. 

Carbon. Symbol, C. Atomic Weight, 12. 

170. History and Occurrence. Carbon, in most of its 
forms, has been known since the very earliest times, 
although its allotropic modifications were not recognized 
until the end of the eighteenth century, and its relations 
to organic chemistry were not thoroughly understood 
until some years later. 

Carbon is a very abundant element. It occurs in nature 
free as diamond, graphite, and the different varieties of 
mineral coal. Combined with other elements, principally 
hydrogen, oxygen, and nitrogen, it occurs in a greater 
number of. compounds than any other known element, be- 



140 LECTURE NOTES 

ing found in all animal and vegetable substances, and their 
derivatives. Combined with oxygen alone, it occurs as 
carbon dioxid in the atmosphere, and combined with oxy- 
gen and the metals, it forms the carbonates, which con- 
stitute quite a portion of the earth's crust. The carbon- 
ates are also found in small quantities in nearly all natural 
waters. 

171. The Allotropic Forms of Carbon. There are three 
allotropic forms of carbon, viz : diamond, graphite and 
amorphous carbon. There are many varieties of the last 
form. In each of these forms carbon is an infusible, non- 
volatile, insoluble solid, having neither odor nor taste. 
The different forms differ very much in physical properties, 
such as color, hardness, specific gravity, luster, and in 
most other properties. When strongly heated in oxygen 
they all form carbon dioxid. 

Diamond occurs in nature in crystals which belong to 
the isometric system. It occurs in small quantities, and 
few localities. The crystals are usually colorless, or 
slightly yellow, are transparent, and have an adamantine 
luster. Diamond has a specific gravity of 3.5 and is the 
hardest substance known. 

Graphite also occurs in nature quite widely distributed 
and in considerable quantities. It usually occurs incom- 
pact foliated masses, although it is occasionally found in 
six-sided tabular crystals belonging to the hexagonal 
system. It is a soft, shiny, grayish-black substance, is 
smooth and slippery to the touch, and entirely opaque. 
It has a specific gravity of about 2.2, and is so soft that 
it leaves a black streak on paper. 

Amorphous carbon occurs in nature in the various 
forms of mineral coal. It is produced artificially by the 



LECTURE NOTES 141 

incomplete combustion of organic substances. It is never 
found, or produced, perfectly pure, although in some of its 
forms it is very nearly so. The specific gravity of the 
different forms varies from 1.2 to 2.3. They vary also 
very much in hardness, some kinds being quite soft, while 
others are very hard. 

DIAMOND 

172 Occurrence and Properties. The diamond is one 
of the rarest and most valuable of mineral gems and has 
been highly regarded for ages on account of its brilliant 
luster and remarkable hardness. It was at first believed 
that the diamond was a peculiar kind of rock-crystal, but 
in 1777 Bergman proved that the diamond contained no 
silica, and a little later it was clearly proved that it was 
composed of nothing but carbon. 

Diamonds occur mostly in alluvial deposits of gravel, 
sand or clay, in connection with minerals which are com- 
monly found in granitic rocks. They are also found in 
the neighborhood of a kind of micaceous rock called itacol- 
umite. This rock was first noticed in Brazil and is dis- 
tinguished by being quite flexible. They have been found 
in a few places in the original rock, and also in a few 
meteorites. 

The principal diamond localities are British India, Brazil 
and South Africa. Most of the diamonds at the present 
time, come from the South African mines, where they are 
found in the so-called "blue ground." They have occa- 
sionally been found in other localities, but the total amount 
found elsewhere is very small. 

Diamonds occur in more or less modified octahedral 
crystals, which belong to the isometric system. The 



142 LECTURE NOTES 

crystals often have curved faces and edges. They are 
usually colorless or slightly yellow, but are occasionally 
found colored green, blue, red, or black. 

Imperfect diamonds, which cannot be used as gems, are 
called boort, or bort. The boort is used in the form of 
powder for cutting and polishing diamonds. Fragments 
of boort are used for writing upon glass, while the natu- 
rally curved edge of a crystal is used for cutting glass. 
They are also used for making the so called diamond drills 
which are employed in rock-boring. An impure black 
variety of diamond is called carbonado, and is also used 
for the same purposes. 

The natural brilliancy of the diamond is greatly in- 
creased by cutting and polishing. This is done by holding 
the surface of the stone against a rapidly revolving metal 
wheel, covered with diamond dust and oil. 

If the diamond is strongly heated in the air, or in 
oxygen, it will burn, forming carbon dioxid. It is en- 
tirely unaffected by most chemical reagents. It combines 
with sulfur when heated to 1000°, and, when fused with 
sodium or potassium carbonate disappears, forming car- 
bon monoxid. 

The unit of weight for the diamond is the carat. This 
is a very old weight, and is equal to 0.205 grams, or 3.16 
grains. The value of the diamond increases very rapidly 
with its size, the increase being approximately as the 
square of its weight. 

Among the famous large diamonds may be mentioned 

the Koh-i-noor, (106 carats), the Regent, or Pitt, 

(137 carats), the Tuscan or Austrian, (139^ carats), 

and the Orloff, (195 carats), all found in India ; the 



LECTURE NOTES 143 

Star of the South, (125 carats), found in Brazil ; the 
Victoria, (180 carats), the Tiffany, (125 carats), and 
others, found in South Africa. Of colored stones the 
most famous are the Hope, a blue diamond, (44 carats), 
the Dresden green diamond, (40 carats), and the red 
diamond of Czar Paul, (10 carats). The Tuscan and the 
Tiffany are light yellow. 

173. Preparation of the Diamond. A great many at- 
tempts have been made to produce diamonds artificially, 
all of which have failed until quite recently. Nothing 
was actually known of the method by which they were 
produced in nature, until their discovery in meteoric iron, 
which seemed to indicate that they had been formed by 
crystallization from molten iron, and under high pressure. 
This process was successfully carried out by Moissan, 
in 1893, who succeeded in preparing both diamond and 
carbonado artificially. 

In order to prepare artificial diamond some pure sugar 
charcoal is compressed in a cylinder of soft iron 
and tightly closed with a plug of the same metal. This 
is then placed in a crucible containing molten wrought 
iron, and heated in an electric furnace to a very high 
temperature. The crucible is then quickly withdrawn, 
and the iron cooled in melted lead. Iron, like water, 
expands slightly when it solidifies, and so when the out- 
side portion solidifies, the interior as it cools, is subjected 
to enormous pressure. After it has cooled, the whole 
mass is treated with hydrochloric acid, which dissolves the 
iron, when the residue is found to consist of graphite and 
diamond, a portion of the latter being colorless and trans- 
parent, and showing all the properties of the natural 
diamond. 



144 LECTURE NOTES 

GRAPHITE 

174. Occurrence, Preparation, and Properties. Graph- 
ite occurs quite widely distributed. It is usually found 
in compact foliated masses, or in hexagonal plates, the 
principal localities being Cumberland, in England, Ceylon, 
Siberia, and several localities in the United States. 

Graphite can be formed artificially, by fusing carbon 
with iron. A small amount of the carbon dissolves in 
the molten iron, and crystallizes upon cooling, forming 
graphite, which can be obtained by dissolving the iron in 
hydrochloric acid. 

Graphite is a grayish-black solid, entirely opaque, and 
unctuous to the touch. It conducts electricity, and is so 
soft that it leaves a black streak on paper. It was at first 
supposed to contain lead, and so was called plumbago or 
black lead. It does not combine with oxygen except at a 
very high temperature, being quite as difficult to burn as 
the diamond. Its principal uses are in the manufacture of 
pencils, crucibles and stove polish. Most graphite con- 
tains a small percentage of hydrogen, showing that it is 
probably of vegetable origin. 

Graphite in its chemical relations differs considerably 
from the other forms of carbon. If it is treated with 
fuming nitric acid and potassium chlorate, it is changed 
into what is called graphitic acid, C n H 4 0.. 

AMORPHOUS CARBON 

175. General Properties and Varieties. A great num- 
ber of varieties of amorphous carbon exist. They are 
grayish-black or black and opaque, and differ considerably 
in their physical properties. Amorphous carbon occurs 



LECTURE NOTES 145 

in nature in enormous deposits under the general name of 
coal, of which there are many sub-varieties. The other 
varieties of amorphous carbon are classified according to 
their origin and properties as charcoal, coke, gas-carbon, 
and lamp black. 

All forms of amorphous carbon are of vegetable or 
animal origin, and differ very much in purity, according to 
their source and the completeness of the decomposition 
of the substances from which they were obtained. 

176. Coal. When vegetable matter decays in the ab- 
sence of air, in the earth, or under water, a number of 
gaseous substances are given off, and there remains a 
substance rich in carbon which we call coal. There are 
many sub-varieties of coal, all of which may be rather 
broadly classified under two divisions, hard coal or an- 
thracite coal, and soft coal or bituminous coal. These two 
divisions depend upon the completeness of the decomposi- 
tion of the vegetable material from which the coal was 
derived. 

Anthracite coal has undergone the most complete 
decomposition and so has the largest percentage of car- 
bon. It is a black, brittle solid, having a bright, almost 
metallic luster, is often iridescent and is quite hard. It 
has a specific gravity of 1.3 to 1.8 and contains from 90 
to 91 per cent, of carbon. The best kinds burn with no 
smoke and very little flame. 

Bituminous coal includes a number of varieties, all of 
which burn with a smoky flame, and, on distillation, yield 
a number of volatile hydrocarbons, together with tar, 
or bitumen. 

Bituminous coal is of two kinds, caking and non- 



146 LECTURE NOTES 

caking. Caking coal softens and partially fuses in the 
fire, whi]e the non-caking coal does not. 

Cannel coal differs from ordinary bituminous coal, 
principally in texture. It yields a large amount of vola- 
tile products and so is especially adapted for gas-making. 

There are a number of varieties of coal belonging to a 
later geological period than those described. These con- 
tain from 60 to 80 per cent, of carbon, and are used for 
fuel. The principal ones are boghead coal, brown coal, 
and peat, or turf. 

177. Charcoal. Common charcoal is a substance which 
is made by the imperfect combustion of wood. Large piles 
of wood are covered with sods and earth, and allowed to 
burn, with an insufficient supply of air, until all the vola- 
tile products are driven off. 

Charcoal is a black lustrous solid. It has neither odor 
nor taste and burns without smoke or flame. It is used 
principally for the reduction of ores, and for fuel. 

Charcoal is very porous, and has the property of ab- 
sorbing gases, especially ammonia and hydrogen sulfid, 
to a large extent. It is, on this account, used in filters 
for purifying drinking water, and as a disinfectant and 
deodorizer. 

The purest form of charcoal is that made by heating 
sugar and igniting the product in a stream of chlorin. 

The specific gravity of charcoal is about 1.5. Ordinary 
charcoal floats on water because its pores are filled with 
air which makes its specific gravity, under such conditions, 
only about 0.2. If a piece of charcoal is boiled in water, 
the air will be driven out of the pores and then the char- 
coal will be seen to be heavier than water. 



LECTURE NOTES 147 

Animal charcoal is prepared by heating all kinds of 
animal refuse, including the bones, in iron retorts. It 
consists of about 10 per cent, of carbon, the rest being 
mostly calcium phosphate. It possesses absorptive pow- 
ers for coloring matter and gases in a much higher degree 
than common charcoal. It is principally used in sugar 
refineries for clarifying the syrup. 

178. Coke. Coke bears the same relation to coal as 
charcoal does to wood. In the manufacture of illuminating 
gas, bituminous coal is heated in retorts, until all the 
volatile products have been driven off, when coke remains. 
It is also made in coke ovens, by burning coal to a certain 
point, and then stopping the combustion. It is a gray, 
lustrous, porous solid, which takes fire at a much higher 
temperature than common coal, and burning, gives a 
very high temperature. It is used in iron smelting. 

179. Gas-Carbon. Gas-carbon is found as a deposit in 
the upper portion of the gas retorts in the manufacture 
of coal gas. It is also formed when ethylene, C 2 H 4 , is 
made to pass through a red hot tube. It is an iron -gray 
porous mass, which is very hard, and conducts electricity. 
It is a very pure form of carbon and is used for making 
the carbon plates in batteries, and for the candles used 
in electric lighting. 

180. Lampblack or Soot. When any compound rich in 
carbon, is burned with an insufficient supply of air, some 
of the carbon passes off, forming what we call lampblack 
or soot. Resin, turpentine, or crude petroleum may be 
burned for this purpose. It is purified by heating in a 
closed vessel, although after this it still retains hydrogen, 
which can only be removed by long continued heating in 



148 LECTURE NOTES 

a stream of chlorin. This is one of the purest forms of 
amorphous carbon, and is used in making India ink, and 
as a pigment in calico printing. Common lampblack is 
used for black paint, and for printer's ink. 

COMPOUNDS OF CARBON AND HYDROGEN 

181. The Hydrocarbons. Carbon forms a very large 
number of compounds with hydrogen, which are known 
as the hydrocarbons. 

The carbon atoms possess in a marked degree, the prop- 
erty of uniting among themselves ; that is, two carbon 
atoms will partially satisfy their valence between them- 
selves, leaving the rest to be satisfied by other atoms. 

This gives rise to a number of groups, or series of 
compounds, the members of which are related to each 
other in the simplest manner. Each compound in each 
series, differs from the preceding compound by the incre- 
ment CH 2 . There are three of these series that are of 
considerable importance. These may be expressed by the 
general formulas, 

C n H. 2nf 2> C n H 2n . and C n H 2n _ 2 . 

Since all of these compounds belong to organic chemis- 
try only the simplest one of each of the above series 
will be considered here. These are, 

Methane, CH , 

Ethylene, C 2 H 4 , and 

Acetylene, C 2 H 2 . 

Methane, or Marsh Gas. Molecular Formula, CH 4 , 
Density, 8. 

182. History and Occurrence. Methane has been known 
since very early times. It was not distinguished from 



LECTURE NOTES 149 

hydrogen until the latter part of the eighteenth century, 
when Vol ta showed that in burning, it required four times 
as much oxygen as did hydrogen, and that it produced 
carbon dioxid, which the hydrogen did not. To distin- 
guish it from ethylene, which was the only other hydro- 
carbon known at this time, it was called light carburetted 
hydrogen, ethylene being relatively much heavier. 

It occurs free in nature, being formed by the decompo- 
sition of vegetable matter under water. The bubbles of 
gas which rise when a stagnant pool is stirred, consist 
largely of methane, hence the name marsh gas. It is 
also formed by the slow decomposition of bituminous coal, 
forming what the coal miners call fire damp. It is the 
principal ingredient in natural gas. 

183. Preparation. Methane is best prepared by heating 
a mixture of sodium acetate and sodium hydroxid, the 
latter being in excess, thus : 

NaC 2 H 3 2 -f- NaOH = Na 2 C0 3 -f- CH 4 . 

Thus obtained, it always contains some hydrogen, and 
ethylene. If required pure, it may be obtained by decom- 
posing zinc methyl, Zn(CH 3 , i 2 , by means of water, thus : 

Zn(CH 3 ) 2 + 2 H 2 = Zn(OH), + 2 CH 4 . 

It is also produced by the decomposition of many 
organic compounds. 

184. Properties. Methane is a colorless, odorless, and 
tasteless gas. Next to hydrogen, it is the lightest gas 
known. It is combustible, burning with a slightly lumin- 
ous flame, but does not support combustion. When mixed 
with two volumes of oxygen, or ten volumes of air, it is 
very explosive. It is this latter mixture which is the 



150 LECTURE NOTES 

cause of the terrible explosions which sometimes happen 
in coal mines. When mixed with chlorin and exposed to 
direct sunlight, the hydrogen will be partly or entirely 
replaced by the chlorin. It can be condensed to a liquid 
under a very great pressure at a very low temperature. 

Ethylene. Molecular Formula, C 2 H 4 . Density 14. 

185. History, Occurrence, and Preparation. Ethylene 
was probably discovered by Becher, in the middle of the 
17th contury, although its true composition and proper- 
ties were not described until more than a century later. 

It is found in small quantities in natural gas, and is the 
most important constituent of coal gas, which is made by 
the destructive distillation of coal. 

It is best prepared by the action of a large excess of 
sulfuric acid, upon common, or ethyl alcohol. In this 
action the acid abstracts water from the alcohol, thus : 

C 2 H 5 OH = C 2 H 4 + H 2 0. 

186. Properties. Ethylene is a colorless gas, having a 
pleasant, ethereal odor. It is a trifle lighter than air, its 
specific gravity being 0. 978. It is relatively much heavier 
than methane, and so, to distinguish it, was called heavy 
carburetted hydrogen. It is a combustible gas, burning 
with a highly luminous flame, and forming carbon dioxid 
and water. When mixed with three volumes of oxygen 
and ignited, it is very explosive. 

The most characteristic property of ethylene, is its 
power of combining directly with chlorin to form an oily 
liquid called ethylene dichlorid, C 2 H 4 C1, or Dutch liquid, 
it having been first observed by four Dutch chemists. 



LECTURE NOTES 151 

From this property it received the name olefiant gas, by 
which it was for a long time known. 

Ethylene can be condensed to a liquid, its critical tem- 
perature being 10°, and its critical pressure, 51 atmos- 
pheres. Liquid ethylene boils at — 103°, and by rapid 
evaporation a temperature of — 140° can be obtained. It 
is, therefore, a very useful refrigerating substance, when 
very low temperatures are required. 

Acetylene. Molecular Formula, C 2 H 2 . Density, 13. 

187. Preparation. Acetylene was discovered in 1836. 
It is formed synthetically when a powerful electric current 
is passed through carbon electrodes in an atmosphere of 
hydrogen. It is also formed by the incomplete combus- 
tion of many hydrocarbons. Thus when the gas in a 
Bunsen lamp burns at the bottom of the tube, acetylene 
is formed. It can be conveniently prepared by heating 
ethylene dibromid, with a concentrated solution of potas- 
sium hydroxid in alcohol, thus : 

C H Br + 2 KOH = 2 KBr -f 2 HO if C,H, 

Acetylene is best prepared by decomposing calcium 
carbid, CaC 2 , by means of water, thus : 

CaC 2 + 2 H 2 = Ca(OH) 2 -f C 2 H 2 . 

188. Properties. Acetylene is a colorless gas. It has 
a peculiar and rather disagreeable odor, and is the only 
hydrocarbon which has been formed by direct union of 
the elements. It is soluble in its own volume of water 
and is much more soluble in alcohol. It is combustible, 
and burns with an intensely luminous, but smoky 
flame. It unites directly with the halogens forming 
addition products. It is quite poisonous when breathed, 
combining with the haemoglobin in the blood. 



152 LECTURE NOTES 

Acetylene forms very characteristic compounds with 
certain metals and metallic salts. Thus, if the gas is 
made to pass through an ammoniacal solution of cuprous 
chlorid, Cu 2 Cl 2 , a dark red precipitate is formed. This has 
the composition Cu ? C 2 , aud is very explosive. It forms 
an analogous compound with silver. 

Acetylene has recently become an important commer- 
cial product, being used quite extensively as an illumina- 
ting gas. It burns with an intense white light, and has 
an illuminating power of about fifteen times that of coal- 
gas. 

189. Coal Gas. It was early known that when coal 
was heated, a combustible gas was produced ; but it was 
not until 1792 that Murdock made gas illumination a 
practical success. 

Coal gas is ordinarily prepared, by distilling bituminous 
coal at a high temperature. Other kinds of coal, as well 
as wood, and petroleum, are sometimes employed, the 
process being much the same. 

The coal is heated in iron retorts, and the volatile pro- 
ducts pass first into a large receiver, where tar, and some 
water and ammonia are collected. The gas then passes 
through what are called the scrubbers, where the rest of 
these impurities are collected, then through a purifier, 
where sulfur compounds, and carbon dioxid are absorbed, 
and finally into the gasometer, from whence it is distrib- 
uted in pipes. 

Coal gas is a mixture of several gaseous products, and 
varies somewhat in composition according to its source, 
and the temperature at which it is produced. The mix- 
ture may be divided into three classes ; those which burn 



LECTURE NOTES 153 

with a luminous flame, called illuminants ; those which 
burn with a non-luminous, or only slightly luminous flame, 
called diluents ; and impurities, which have not been, or 
cannot be removed. The principal illuminants are ethy- 
lene, acetylene, and other hydrocarbons of the same series. 
The principal diluents are hydrogen, methane, and carbon 
monoxid. The impurities are nitrogen, carbon dioxid 
and hydrogen sulfid. Illuminating gas contains from 4 
to 11 per cent, of illuminants, from 85 to 90 per cent, of 
diluents, and from 3 to 6 per cent, of impurities. 

The illuminating power of coal-gas, is determined by 
an instrument called a photometer, in which the amount 
of light given by the gas, when it burns at the rate of 
five cubic feet in an hour, is compared with the light of a 
standard candle, which burns 7.79 grams (120 grains) of 
spermaceti per hour. Good coal-gas should have from 16 
to 20 candle power, or even more. 

190. Combustion. When two substances combine chem- 
ically, it frequently happens that the action is accompanied 
by light and heat. Such a phenomenon we call combus- 
tion. We may, therefore, define combustion, as the 
chemical union of tioo sicbstances which combine with 
sufficient energy to produce light and heat. 

Sometimes a chemical change takes place, which is 
exactly analogous to combustion, except there is no light. 
and little heat. Such a change is seen in the decay of 
wood, and the rusting of metals. These changes are 
distinguished as slow combustion. A similar change takes 
place in the lungs, where the blood is purified by oxidation. 

In all cases where combustion takes place, especially 
where the combustion is rapid, we are accustomed to 



154 LECTURE NOTES 

regard one of the substances as the combustible, and the 
other as the supporter of combustion. One of the sub- 
stances in a combustion is usually a gas, which entirely 
surrounds the other. The surrounding medium is the 
supporter of combustion. Thus when a jet of coal-gas 
burns in the atmosphere, or a piece of sulfur burns in 
oxygen, the atmosphere and the oxygen are the support- 
ers of combustion, and the coal-gas and the sulfur are the 
combustibles. 

We happen to live in an atmosphere, the active ingredi- 
ent of which is oxygen ; and so substances which combine 
with oxygen are said to be combustible, while oxygen, and 
those substances which act like oxygen, are called the 
supporters of combustion. In any combustion especially 
if the two combining substances are gases, it is obvious 
that the terms combustible and supporter of combustion 
are entirely relative ; for while coal-gas will burn in the 
air or oxygen, oxygen or air will burn in coal-gas. 

When combustion takes place so rapidly as to be at- 
tended with more or less noise, we call it explosion. 

191. Temperature of Ignition. In order that combus- 
tion may take place, the combustible substance, must be 
raised to a certain temperature. This is called the tem- 
perature of ignition, and varies widely with different 
substances. A few compounds, such as liquid phosphin, 
ignite at the ordinary temperature, when exposed to the 
air. Phosphorus ignites at about 60°, carbon disulfid at 
149°, and sulfur at 260°, while coal-gas will not ignite at 
red heat, and nitrogen only in the electric arc, the highest 
known temperature. 

Slow combustion takes place at a much lower tempera- 
ture. By mechanical division, the exposed surface is 



LECTURE NOTES 155 

increased to such an extent, that the slow combustion, 
beginning at the ordinary temperature, often raises the 
temperature of the substance to the point of ignition. 
This is called spontaneous combustion, and often occurs in 
heaps of old refuse, especially of oily rags, and is a 
common cause of fires. 

If a piece of wire gauze is held over a gas jet, and the 
gas ignited above the gauze, the latter can be raised to 
quite a distance above the jet, and yet the gas will not 
take fire below the gauze. This is due to the fact, that 
the wire in the gauze conducts away the heat so rapidly, 
that the temperature of the gas below, never reaches the 
point of ignition. This principle has been utilized in the 
construction of the safety lamp, invented by Sir Hum- 
phrey Davy, for use in coal mines. 

192. The Nature of Flame. The combustion of sub- 
stances, may, or may not, be accompanied by flame. A 
non-volatile substance, like carbon, burns with a bright 
glowing, but without flame. A combustible gas, or a 
solid or liquid, which at the temperature of combustion 
produces a combustible gas, burns with a flame. We may, 
therefore, define flame, as gas raised by combustion to 
incandescence. 

The flame of a burning candle, or of coal-gas, may be 
divided into four parts, — the dark central portion of un- 
burned gas, the blue region at the base of the flame, the 
luminous portion, or area of incomplete combustion, and 
the thin outer, or non-luminous portion, in which com- 
plete combustion has taken place. 

The luminosity of a flame, depends principally upon the 
presence, in the flame, of solid particles, which are not 
volatile, and which are heated to whiteness by the process 



156 LECTURE NOTES 

of combustion. It also depends upon the pressure. In 
the ordinary gas, or candle flame, these solid particles are 
unburned carbon. A non-luminous flame, like that of 
burning hydrogen, may be made highly luminous, by 
burning it under high pressure, or by introducing some 
solid, like a piece of platinum wire, into the flame. 

If a vessel is heated in a luminous gas flame, soot is 
deposited. To avoid this, we use, for heating purposes, 
a lamp so constructed as to insure complete combustion. 
Such a lamp was invented by Bunsen, and is known as 
the Bunsen lamp. 

Since hot carbon combines with oxygen with great 
energy, and so is reducing in its action, the luminous gas 
flame, which contains unburned carbon, is a reducing 
flame. The non-luminous gas flame, such as that produced 
by the Bunsen lamp, is, on the contrary, an oxidizing 
flame, because it contains an excess of heated air or 
oxygen. 

193. The Halogen Compounds of Carbon. Carbon does 

not unite with the halogens directly, but compounds can 
be formed by the action of the halogens on certain hydro- 
carbons. In this action, the halogen replaces the hydro- 
gen, and substitution products are formed, thus : 

CH 4 + Cl 2 = CH 3 C1 + HC1. 

By continuing the action, and raising the temperature, 
all the hydrogen may be replaced. An intermediate 
compound, is the one known as chloroform, CHC1 3 , which 
is much used as a solvent for fats, rubber, iodin, etc. 

A more complete description of these compounds, be- 
longs to organic chemistry. 



LECTURE NOTES 157 

COMPOUNDS OF CARBON WITH OXYGEN 
Carbon forms but two oxids. These are : 
Carbon monoxid, CO, and 
Carbon dioxid, C0 2 . 

Both compounds are gases, and both are quite important. 
Carbon Monoxid. Molecular Formula, CO. Density, 14. 

194. Occurrence and Preparation. Carbon monoxid 
does not occur free in nature, but is produced in large 
quantities by the combustion of carbon in an insufficient 
supply of air. It is the burning of this gas that causes the 
blue flames which may be seen whenever coal is burning. 

Carbon monoxid is best prepared by heating oxalic 
acid, with concentrated sulfuric acid. This removes water, 
from the oxalic acid, and produces equal volumes of car- 
bon monoxid, and carbon dioxid, thus : 

H 2 C 2 4 = H 2 + C0 2 + CO. 

The gases are made to pass through a solution of sodium 
hydroxid, which dissolves the carbon dioxid. 

Carbon monoxid is formed by the incomplete combustion 
of carbon, and also by passing carbon dioxid over red hot 
charcoal, thus : 

C0 2 + C = 2 CO. 

195. Properties. Carbon monoxid is a colorless, taste- 
less, gas, having a slight, but oppressive odor. It is only 
very slightly soluble in water. It is combustible, and 
burns with a bright blue flame forming carbon dioxid. 
It is extremely poisonous, and combines with the haemo- 
globin in the blood. When present in the air, even in 
small quantities, it produces a severe headache, and in 
larger quantities it produces insensibility, and death. 



158 LECTURE NOTES 

The carbon atom in carbon raonoxid is bivalent, as 
shown by the fact that the molecular formula is CO and 
not C 2 2 . It is, therefore, what is sometimes called an 
unsaturated molecule. It combines directly with oxygen 
and chlorin, and is, therefore, a strong reducing agent. 
It combines directly with some of the metals to form a 
peculiar class of compounds, some of which are very 
explosive. 

Carbon Dioxid. Molecular Formula, C0 2 . Density, 22. 

196. History and Occurrence. Up to the middle of the 
16th century, all gases were supposed to be alike, and 
were called air. Carbon dioxid was the first gas to be 
distinguished from common air. At first, owing to its 
source, or properties, it was called by various names, 
such as chalky air, fixed air, and later, carbonic acid. 

It occurs free in the atmosphere, of which it constitutes 
about 0.03 per cent., being formed by the oxidation of all 
organic substances. It is found in soils, and in many 
mineral springs, and not infrequently collects in deep 
wells, cellars, and coal pits, where it is known as, choke 
damp. 

Combined with oxygen and the metals, it forms a vast 
series of compounds called carbonates, of which the most 
important one is calcium carbonate, CaCO a , which exists 
under the different names of calcite, marble, chalk and 
limestone. Among the more important carbonates are 
magnesite, MgCO a ; dolomite, (CaMg)C0 3 ; and siderite, 
FeC0 3 . 

197. Preparation. Carbon dioxid is most conveniently 
prepared, by the action of some acid upon a carbonate. 



LECTURE NOTES 159 

When marble and hydrochloric acid are used, the action 
is as follows : 

CaC0 3 -f 2 HC1 = CaCl 2 + H 2 -f C0 2 . 

It is also produced by heating limestone, in the prepa- 
ration of lime, thus : 

CaC0 3 =CaO + C0 2 . 

It is formed whenever wood, coal, or any organic sub- 
stance, is burned in the air, or in oxygen ; or, when an 
organic compound decays ; or, by the processes of fer- 
mentation and respiration. 

198. Properties. Carbon dioxid is a colorless gas, hav- 
ing a slightly pungent odor, and an acid taste. It is 
neither combustible nor a supporter of combustion. It is 
much heavier than air, its specific gravity being 1.52. 

It is not poisonous but will destroy life by suffocation. 
Air containing as much as 5 per cent, of carbon dioxid 
may be breathed for a very short time without serious 
results, although air containing only 0.2 per cent., might 
prove dangerous to life, if breathed continuously. 

Its presence may be shown, by the white precipitate of 
calcium carbonate, CaC0 3 , which is formed when it is 
brought into contact with lime water, thus : 

Ca(OH) 2 + C0 2 = CaC0 3 + H 2 0. 

At a temperature of 15°, water, dissolves its own volume 
of this gas. If the pressure be increased, the solubility 
is increased in the same ratio, so that under ten atmos- 
pheres pressure, water dissolves ten times as much of the 
gas as under the ordinary pressure. There is a weak 
chemical union between the carbon dioxid and the water, 
forming carbon acid, H 2 CO a . This is an extremely unsta- 



160 LECTURE NOTES 

ble compound, but its salts, the carbonates, are very- 
stable. 

Although carbon dioxid is constantly being formed in 
the atmosphere, the quantity is never much increased. 
This is because it is the principal plant food. The plants 
absorb carbon dioxid, which is decomposed in the cells of 
the plant leaf, setting free oxygen. This action is much 
increased by sunlight. 

199. Liquid and Solid Carbon Dioxid. Carbon dioxid 
can be condensed to a liquid both by cold and pressure. 
It was first liquefied by Faraday, in 1823. Reintroduced 
into separate parts of a bent tube, some sulfuric acid, 
and some ammonium carbonate, after which the tube was 
hermetically sealed. The acid and the carbonate were 
then carefully brought together, which formed carbon 
dioxid, and the pressure of the gas soon became sufficient 
to cause it to liquefy. Liquid carbon dioxid is now manu- 
factured on a large scale, by pumping the gas into heavy 
steel cylinders, with powerful compression pumps, until 
it condenses. 

Liquid carbon dioxid is a colorless, and extremely 
mobile liquid. It is only very slightly soluble in water, 
upon which it floats without mixing. Under the ordinary 
atmospheric pressure it boils at — 78°. 

If liquid carbon dioxid is heated, it expands more rap- 
idly than a gas. It has, therefore, the largest coefficient 
of expansion of any known substance. 

The critical temperature of carbon dioxid is 32° and its 
critical pressure 73 atmospheres. At 0° it liquefies under 
about 38 atmospheres pressure. 



LECTURE NOTES 161 

If liquid carbon dioxid is allowed to escape into the 
atmosphere through a narrow opening, the absorption of 
heat, due to its rapid evaporation, is so great, that a 
portion of the liquid solidifies. The solid carbon dioxid 
is easily obtained in large quantities, by allowing a jet of 
the liquid to pass into a proper receptacle, such as a stout 
bag. 

Solid carbon dioxid is a soft white substance which 
much resembles snow. When exposed to the air it does 
not melt, but passes directly into the gaseous state. It 
can be melted, however, this change taking place at — 65°, 
a remarkable fact, since the liquid boils at a temperature 
considerably below this. It dissolves readily in ether, 
and if this mixture is made to evaporate rapidly, a tem- 
perature of — 110° can be obtained. 

If solid carbon dioxid is pressed into cylindrical 
wooden moulds, the cylinders thus formed have a den- 
sity of about 1.2, and can be preserved for several hours. 

200. Carbonic Acid, H 2 C0 3 . The number of compounds 
of carbon, hydrogen, and oxygen, is very great. They 
include not only a large number of acids, but also a large 
number of entirely different compounds, such as the alco- 
hols, ethers, sugars, etc. These all belong to that division 
of chemistry known as organic chemistry. Only one of 
these compounds, viz., carbonic acid, will be studied here. 

Carbonic acid is formed when carbon dioxid is dissolved 
in water. It exists only in a weak aqueous solution, 
which decomposes as soon as an attempt is made to con- 
centrate it, forming carbon dioxid and water. In this 
solution it acts like a true acid, reddening litmus paper, 
aud combining with the bases to form salts. It is a 
dibasic acid and so forms both normal and acid salts. 



162 LECTURE NOTES 

These salts, which are called carbonates, are quite sta- 
ble, and are among the most important of all chemical 
compounds. 

201. The Carbonates. The carbonates form a very 
important class of natural compounds. The most impor- 
tant of these is calcium carbonate or calcite. This 
substance occurs very widely distributed and often in 
enormous masses. When pure it forms beautiful crystals 
which belong to the hexagonal system, and of which there 
are hundreds of different forms known. The more or less 
impure forms are called marble, chalk, coral, limestone, 
etc. In addition to this compound, carbonates of a large 
number of other metals are found in nature. 

The carbonates can be formed by leading carbon dioxid 
into a solution of any hydroxid in water, thus : 

2 NaOH + C0 2 = Na 2 C0 3 + H,0. 

Carbonates of those metals whose hydroxids are insolu- 
ble in water, can be formed by the action of a soluble 
carbonate upon a solution of any of the metallic salts. 
This very often results in the formation of basic car- 
bonates. 

The acid carbonates are not very numerous. They can 
generally be prepared by the continued action of carbon 
dioxid upon the normal carbonate, either dissolved in 
water or held in suspension by water. Many of these 
decompose by boiling, forming the normal carbonates. 

All carbonates are decomposed by the common acids 
with effervescence, caused by the giving off of carbon 
dioxid. 

202. Carbon Disulfid, CS 2 . Carbon disulfid was acci- 
dentally discovered in 1796, while heating pyrites with 



LECTURE NOTES 163 

charcoal. It does not occur free in nature, but is easily 
formed synthetically, by leading the vapor of sulfur over 
hot charcoal in a suitable retort. 

Carbon disulfid is a colorless, volatile, strongly refract- 
ing liquid. When pure, it has a pleasant, ethereal odor. 
The commercial substance generally contains some sulfur 
and sulfur compounds, which impart to it a very disa- 
greeable odor. It has a specific gravity of 1.29, and 
boils at 46°. Its vapor is very inflammable, and takes 
fire at 149°, burning with a blue flame. When mixed 
with nitric oxid, it burns with great brilliancy. 

Carbon disulfid is much used as a solvent for the fats 
and oils, for rubber, phosphorus and sulfur, and for ex- 
tracting certain essential oils. 

Carbon disulfid is a powerful poison. Its vapor is 
especially fatal in its effect upon small animals and insects, 
and it is therefore used as an insecticide. Its vapor when 
breathed, acts upon the nervous system and is dangerous 
even to the larger animals, and to man, if breathed con- 
tinually, even in small quantities. 

203. Cyanogen Gas, (CN) 2 . Cyanogen is a compound 
radical, which, in its chemical properties closely resem- 
bles the elements of the halogen group. Cyanogen gas 
bears the same relation to the compound radical cyanogen, 
that the chlorin molecule does to the atom of chlorin. 

Cyanogen gas is best prepared by heating mercuric 
cyanid, thus : 

Hg(CN), = Hg + (CN)„. 

It is a colorless gas, and has a peculiar pungent odor, 
somewhat like that of peach kernels. It is combustible, 



164 LECTURE NOTES 

burning with a characteristic reddish-violet flame, and is 
exceedingly poisonous. 

Cyanogen forms one compound with hydrogen. This 
is analogous to the halogen acids, and is called hydrocy- 
anic acid. Its formula is HON. Hydrocyanic acid, which 
is also known commercially as prussic acid, is, when pure, 
one of the most powerful and intensely active poisons 
known, 3.5 milligrams (about -^ of a grain) being usually 
a fatal dose. The salts of this acid, which may also be 
regarded as compounds of cyanogen with the metals, are 
called cyanids. In their chemical properties they much 
resemble the corresponding halogen salts. They are all 
very poisonous. 

Cyanogen forms a large number of compounds and 
derivatives, most of which belong to organic chemistry. 

204. The Double Cyanids. The cyanids of the so-called 
alkali metals, sodium and potassium, and a few others, 
are soluble in water, all the rest being insoluble in water. 
These insoluble cyanids, however, dissolve in the alkaline 
cyanids, forming soluble double cyanids, that is, com- 
pounds in which there are two metals. The most of these 
are decomposed by dilute acids, forming again the insol- 
uble cyanid, and liberating hydrocyanic acid. 

A few of the metals, particularly iron, form with the 
alkaline cyanids double compounds, which do not liberate 
hydrocyanic acid, when acted upon by dilute acids. Thus 
if a solution of ferrous sulfate, FeS0 4 is treated with an 
excess of potassium cyanid, KCN, there is formed the sol- 
uble double cyanid of iron and potassium, Fe(CN) 2 , 4 KCN. 
If this is acted upon by dilute hydrochloric acid, a white 
crystalline compound is formed, which has the composition 



LECTURE NOTES 165 

H 4 Fe(CN) 6 , and which is known as hydroferrocyanic acid. 
In this compound, which is a strong acid, the iron is a 
part of the acid radical, and the hydrogen is replaceable 
by metals, as is the case with ordinary acids. The salts 
are called ferrocyanids. Corresponding compounds ex- 
ist with ferric iron, and are called ferricyanids. 

The metals manganese, cobalt, platinum and other 
metals like platinum, give analogous series of compounds. 

Silicon, Symbol, Si. Atomic Weight, 28.4 

205. Occurrence. Silicon, next to oxygen, is the most 
abundant element in nature. Over 27 per cent, of the 
solid earth is silicon. 

Compounds containing silicon have been known since 
very early times, and minerals rich in silica were used in 
glass making by the ancients. 

Silicon is never found free in nature, but always in 
combination with oxygen, with which it forms silicon 
dioxid, or silica, SiO z . Silica is very abundant in nature, 
where it occurs as quartz, flint, opal, etc. , and in a more 
or less impure form as sand, and sandstone. 

In combination with oxygen and the metals it forms 
the large and very important class of minerals called the 
silicates, which constitute the greater portion of the 
earth's crust. 

206. Preparation. Silicon was first prepared in an 
impure form, by Berzelius, in 1810. He obtained it by 
fusing together iron, carbon, and silicon dioxid. It can 
be prepared fairly pure, by heating to redness in an iron 
tube, a mixture of potassium silico-fluorid and metallic 
potassium. The action, which is quite violent, is as follows : 

K 2 SiF 6 + K 4 = 6KF-^Si. 



166 LECTURE NOTES 

The potassium fluorid is then dissolved in water, the 
silicon remaining as an insoluble powder. 

If not required absolutely pure, silicon can best be pre- 
pared by G-attermann's method. This consists in heating 
a mixture of powdered sand and magnesium. If only a 
small quantity is required, the operation may be carried 
on in a test tube. An excess of the sand, in the propor- 
tion of about four to one of the magnesium, should be 
used. The action is as follows : 

Si0 2 -f 2 Mg = 2 MgO -f Si 

The action is quite vigorous, and the silicon obtained, 
while not pure, is pure enough for most purposes. 

207. Properties. When prepared as above, silicon is a 
brown amorphous powder, which is insoluble in water, 
and in all mineral acids, except hydrofluoric acid, but 
dissolves in the alkalis forming the alkaline silicates, 
thus : 

Si + 2 KOH + H 2 = K 2 SiO s + 2 H 2 . 

Silicon has a specific gravity of 2.49. When heated 
in the air, or oxygen, it takes fire and burns, forming 
silicon dioxid. 

If amorphous silicon is strongly heated in the absence 
of air, it becomes denser and after a time assumes an 
appearance much like graphite. This is in all probability 
an allotropic form. It oxidizes with great difficulty and 
is insoluble in hydrofluoric acid. 

If amorphous silicon is fused with zinc, it forms dark 
glittering, octahedral crystals, which may be obtained by 
dissolving the zinc in some acid. Crystallized silicon is 
not acted upon by hydrofluoric acid, is oxidized with 



LECTURE NOTES 167 

difficulty in oxygen, and is so hard that it will scratch 
glass. 

208. Silicon Hydrid, or Silico-Methane, SiH 4 . This 
compound is obtained in an impure condition, mixed with 
hydrogen, by the action of hydrochloric acid on magne- 
sium silicid. This latter compound is obtained by heat- 
ing a mixture of powdered sand and magnesium in the 
proportion of two parts sand to three of magnesium. The 
action is quite violent and on a small scale may be carried 
on in a test tube. (See 206). The formation of the sili- 
con hydrid is as follows : 

Mg 2 Si -f 4HC1 = 2 MgCl 2 -f SiH + 

Silicon hydrid, called also silico-methane from its anal- 
ogous compound methane, is a colorless gas. It is 
inflammable, and, when pure, takes fire at a very little 
above the ordinary temperature ; but when mixed with 
hydrogen, as prepared above, it is spontaneously inflam- 
mable, burning with a brightly luminous flame, and giving 
off silicon dioxid. It is decomposed at a red heat into 
amorphous silicon and hydrogen. 

COMPOUNDS OF SILICON WITH THE HALOGENS 

Silicon forms compounds with each of the halogens. 
The most of them are unimportant. 

209. Silicon Tetrachlorid, SiCl 4 . This compound is 
formed by leading dry chlorin over a very strongly heated 
mixture of silicon dioxid, and carbon. The action is as 
follows : 

Si0 2 +.2 C -f 2 Cl 2 = 2 CO -l Si01 4 . 

It is also formed by heating the impure silicon obtained 
by Gattermann's method, to about 300°, and leading over 



168 LECTURE NOTES 

it dry chlorin. [The corresponding bromid and iodid may 
be formed in a similar way.] 

Silicon tetrachlorid is a colorless, volatile liquid, boiling 
at 59°. When brought into contact with water, it at 
once decomposes, forming gelatinous silicic acid, and 
hydrochloric acid. The bromid and iodid act in a similar 
way. 

210. Silicon Tetrafluorid, SiF 4 . This compound is 
formed when sand, or powdered glass, or any compound 
containing; silicon dioxid, is heated with calcium fluorid 
and sulfuric acid, thus : 

Si0 2 -f 2 CaF 2 -f- 2 H 2 S0 4 = SiF 4 + 2 H 2 + 2 CaS0 4 . 

It is this compound which is formed when glass is etched 
with hydrofluoric acid. (See 90.) 

Silicon tetrafluorid is a colorless gas. It has a pungent, 
suffocating odor, somewhat like hydrochloric acid, and 
fumes strongly in the air. It is decomposed by water 
forming silicic acid, and a peculiar compound called hydro- 
fluosilicic acid, thus : 

3 SiF 4 + 4 H 2 = H 4 Si0 4 + 2 H 2 SiF 6 . 

This latter compound is a liquid, having strongly acid 
properties, and forms a well defined series of salts which 
are called silico-fluorids. 

COMPOUND OF SILICON WITH OXYGEN 

Silicon forms but one compound with oxygen. 

Silicon Dioxid, or Silica. Molecular Formula Si0 2 . 

211. Occurrence. Silicon dioxid, which is more com- 
monly called silica, is a very abundant compound. The 



LECTURE NOTES 160 

most important form of silica found in nature is quartz. 
This occurs as hard, usually transparent, well denned 
crystals, which belong to the hexagonal system. 

The purest form of quartz is colorless and transparent, 
and is often called rock crystal. Quartz sometimes occurs 
colored purple or violet, and is then called amethyst. This 
color is supposed to be due to manganese. It is found 
colored smoky-yellow or brown and transparent, to nearly 
black and opaque, the color being due to some compound 
of carbon. This is called smoky quartz. In a massive 
form it is often colored rose, yellow or blue. In an im- 
pure form it occurs as sandstone and sand. 

In addition to the crystalline varieties there are several 
so-called cryptocrystalline varieties of quartz. These are 
often colored by certain metallic oxids, and are known as 
chalcedony, carnelian, agate, onyx, jasper, flint, etc. 

An amorphous form of silica, called opal, is also found 
in nature. 

Silica occurs as a fine powder in large deposits in vari- 
ous parts of G-ermany. This is composed of the shells of 
extinct diatoms and is known as kieseh/iihr or tripolite. 

Silica is also found in the straw of certain cereals, and 
in the feathers of certain birds. 

212. Preparation and Properties. Silica is formed by 
burning silicon in the air, or in oxygen. It can be 
formed by decomposing the alkaline silicates by means of 
an acid, and is easily obtained by heating silicic acid, 
thus : 

H 4 Si0 4 = 2H 2 + Si0 2 . 

Prepared in this way silica is a white amorphous pow- 
der. It is insoluble in water, and in all acids, except by- 



170 LECTURE NOTES 

drofluoric acid, but dissolves easily in the strong alkalis 
forming soluble silicates. 

Silica can be fused in the oxy-hydrogen flame, and can 
be made to boil in an electric furnace. The specific 
gravity of quartz crystals is 2.65, while that of amor- 
phous silica is only 2.2. 

Silica, in its different forms, is extensively used in the 
manufacture of glass, porcelain, and the various kinds of 
pottery. The transparent crystals of quartz, both colored 
and colorless, as well as most of the other varieties, are 
employed for ornamental purposes, and as gems. Agate 
is used for making the agate mortars so much used by 
chemists. 

Tripolite is used extensively in the manufacture of 
dynamite, as an absorbent for the nitro glycerol; also as 
a non-conducting material, for packing steam pipes, and 
for filtering. 

213. Glass. This important commercial product is 
composed chiefly of silicon dioxid. It is made by fusing 
pure sand, or quartz, together with some alkaline carbon- 
ate, and some metallic oxid. These compounds combine 
on heating, forming silicates of the alkali and the me- 
tallic oxid, which, together with a large amount of free 
silica, constitutes glass. The alkaline carbonate may be 
of sodium or potassium while the principal metallic oxids 
used are those of calcium, lead, and iron. 

The larger the proportion of silica, the more infusible 
is the glass, and the less is it acted upon by other sub- 
stances. Potassium glass is more infusible than sodium 
glass. Calcium adds to the hardness as well as to the 
infusibility. Lead makes the glass much softer, more 



LECTURE NOTES 171 

easily fusible and adds greatly to its refracting power 
for light. Iron is only used in the cheaper bottle glass. 

Glass is a hard, brittle substance, which is transparent 
and may be obtained in almost any shade of color by the 
addition of different metallic oxids. 

Porcelain is somewhat analogous to glass. Its princi- 
pal ingredient is kaolin, which is a fairly pure silicate of 
aluminum. This makes a substance which is opaque and 
infusible. 

The different kinds of pottery are similar to porcelain 
but common clay is used in place of the kaolin. 

214. The Silicic Acids. If any silicate is fused with 
sodium or potassium carbonate, a soluble silicate is formed. 
This is easily decomposed by hydrochloric acid, forming 
free silicic acid, a portion of which separates out as a 
gelatinous mass while the rest remains in solution. If 
the solution is sufficiently dilute, all the silicic acid may 
remain in the solution. Silicic acid is also formed when 
any of the halogen compounds of silicon are decomposed 
by water. 

If the silicic acid obtained by any of these methods 
remains in solution, it may be separated from the other 
compounds by the process called dialysis. This is done 
by means of an instrument called a dlalyzer, which con- 
sists of a shallow drum, the bottom of which is a piece of 
parchment paper. In this is placed the mixture to be 
separated, and the whole is then placed on the surface of 
a large volume of water. The crystalline salts and the 
hydrochloric acid diffuse readily through the parchment 
paper, while the silicic acid, is unable to pass through, 
and so remains in the dialyzer in a fairly pure condition. 



1 72 LECTURE NOTES 

Substances which pass readily through a porous mem- 
brane are called crystalloids, while other substances which 
form jellies, and which do not pass through, are called 
colloids. 

There are several silicic acids. The normal acid, HSiO,, 
is the one from which all the others are derived. Meta- 
silicic acid, H 2 Si0 3 , is derived from the normal acid by 
the removal of one molecule of water. The polysilicic acids 
are derived from two or more molecules of the normal 
acid by loss of water. The most important of these acids 
have the composition H 2 Si 2 0. and H 4 Si 3 8 ; some of them 
are very complicated. 

The silicic acids are all jelly-like solids, which, when 
dried, form a white, amorphous powder, and at 100° lose 
all their water, forming silicon dioxid. None of the silicic 
acids possess an absolutely constant composition, and 
none have any value as acids ; but their salts are well 
defined compounds and all are important. 

215. The Silicates. The name silicate indicates a salt 
of some silicic acid, and an extremely large number of 
these compounds are known. They comprise by far the 
greater portion of the earth's crust and some of them are 
extremely complicated. About every possible silicic acid 
is represented in nature by its corresponding silicate. 

Among the more important natural silicates we may 
mention the following, as being typical minerals in their 
respective classes : the orthosilicates include the micas 
and garnets, the metasilicates include the pyroxenes and 
hornblendes, and the polysilicates include the very impor- 
tant group of feldspars. In addition to these general 
classes we find basic silicates and hydrated silicates. In 
the latter class the water sometimes occurs as water of 



LECTURE NOTES 173 

crystallization, and sometimes as chemically combined 
water, or water of constitution. The hydrated silicates 
are decomposed by acids with separation of silicic acid. 

The silicates of the alkali metals are soluble in water. 
These compounds are known as soluble glass. They all 
contain more or less free alkali in the solution, and are 
employed commercially as a fire proof glazing, or paint, 
as a cement for glass, and in the manufacture of artificial 
stone. 

A more exhaustive study of the silicates belongs to the 
science of mineralogy. 

Boron. Symbol, B. Atomic Weight, n. 

216. Occurrence and Preparation. Boron is the only 
non-metal in the group of elements to which it belongs. 
It is never found free in nature, but occurs combined with 
hydrogen and oxygen as boric acid T H 3 BO s , and as 
borates, principally of sodium, magnesium and calcium. 
Its principal compound, borax, has been known since very 
early times, but the element itself was first obtained by 
Davy in 1808. 

Boron may be prepared by igniting boron trioxid with 
sodium or potassium in the absence of air. After the 
action, which is quite violent, has ceased, the fused mass 
is treated with water and a little hydrochloric acid, the 
boron remaining undissolved as a dark brown amorphous 
powder. It can also be prepared from borax by means 
of magnesium. The borax must be fused and perfectly 
dry. The action is much like the one described under 
silicon, only not quite so vigorous, and the temperature 
required is somewhat higher. The fused mass is treated 



174 LECTURE NOTES 

with water and then with hydrochloric acid, which leaves 
the boron fairly pure. 

217. Properties. Boron, prepared as above, is a brown 
amorphous powder. It is infusible even at the tempera- 
ture of the electric arc, but when heated in the air to 
about 700°, it burns. It combines directly with chlorin 
and bromin, and is acted upon by the oxy-acids, forming 
boric oxid and boric acid. It also combines with nitrogen 
directly at a high temperature. 

If amorphous boron is fused with aluminum in a cruci- 
ble, placed in a larger crucible and covered with powdered 
charcoal to prevent oxidation, it crystallizes in the molten 
aluminum and after cooling may be separated from the 
aluminum by dissolving that metal in hydrochloric acid. 
The boron is left in the form of transparent brownish- 
yellow crystals which are called adamantine boron. These 
crystals were at first supposed to be an allotropic form 
of boron ; but investigation shows that they are not 
perfectly pure boron, but that they contain from 2 to 4 
per cent, of carbon. This carbon is believed to be in the 
form of diamond, since those crystals which contain the 
largest amount of carbon have the highest degree of trans- 
parency. According to Hampe. these crystals also con- 
tain aluminum and possess a constant composition, which 
corresponds to the formula B 4g C 2 Al 3 . 

Adamantine boron forms monoclinic crystals which have 
a brilliant lustre and a hardness almost equal to the dia- 
mond. They have a specific gravity of 2.68. When 
heated in the air or in oxygen they oxidize with difficulty, 
and are entirely unaffected by all acids. 

218. The Compounds of Boron. The compounds of 



LECTURE NOTES 175 

boron are for the most part of little but theoretical inter- 
est, and so need but brief mention. In all its compounds 
boron is trivalent. 

Boron forms a compound with hydrogen called boron 
hydrid. If an excess of magnesium is heated with boron 
trioxid they combine with great energy forming a gray 
mass, which contains among other things magnesium 
borid. If this compound is treated with hydrochloric 
acid, boron hydrid is formed, mixed with hydrogen. It 
has never been prepared pure. It is a colorless gas with 
an extremely unpleasant odor, and burns with a bright 
green flame. Its formula has not been determined with 
certainty, but is undoubtedly BH S . 

Boron forms compounds with each of the halogens, 
which in their preparation and properties very closely 
resemble the corresponding compounds of silicon. 

Boron in the amorphous form combines directly with 
nitrogen at a high temperature to form boron nitrid, BN. 
This is a white, amorphous, and perfectly stable compound. 

Boron trioxid, B 2 O g , is the only compound of these two 
elements. It is formed by heating boric acid to redness. 
It is not volatile at a red heat and so many salts are con- 
verted into borates when heated to redness with this 
compound. Most metallic oxids dissolve in fused boron 
trioxid and often impart to it characteristic colors. 

219. The Boric Acids and Borates. Boric acid, also 
called boracic acid, is found free in nature, in the region of 
extinct volcanoes, the most important localities being in 
Tuscany, and in California. It may be prepared by 
decomposing a solution of borax, with sulfuric acid, thus: 

Na 2 B 4 7 + H 2 S0 4 + 5 H,0 = Na,S0 4 + 4 H 3 B0 3 . 



176 LECTURE NOTES 

Boric acid is a white solid, which crystallizes in glisten- 
ing scales. It is somewhat soluble in cold water, and is 
easily soluble in hot water, and in alcohol. If the latter 
solution is ignited, it burns with a characteristic green 
flame. 

When boric acid is heated to 100°, it loses water and 
forms metaboric acid, HB0 2 . At 140° it loses more water 
and forms tetraboric or pyroboric acid, H 2 B 4 O r When 
any of these acids are strongly ignited, all the water is 
driven off, and there remains boron trioxid, B 2 3 . 

There are no salts of normal boric acid known, and the 
'metaborates are mostly unstable, but jfche tetraborates 
are all very stable. 

The most important compound of boron, is sodium 
tetraborate, or borax, Na 2 B + O r This compound finds 
abundant use in the arts as a flux. Most of the metallic 
oxids dissolve in fused borax, and often impart to it 
characteristic colors. It is used in this way in the labo- 
ratory, and serves thus as a means of recognizing many 
of the elements. 



PART THREE 



The Metals 



220. General Characteristics. In studying the classifi- 
cation of the elements (33) it has already been noticed that 
there is no sharp line of separation between the metals and 
the non-metals. In extreme cases the distinction is per- 
fectly clear, and if the elements were always to be studied 
as such, this classification would be comparatively easy; 
but in determining their relations to each other they are 
more commonly studied in compounds, that is, in com- 
bination with other elements. These other elements 
frequently exert such a powerful influence on the first 
element that its nature appears to be changed, and what 
appeared in the elementary state to be a metal, assumes, 
temporarily at least, the properties of a non-metal. This 
makes the division rather arbitrary, and we frequently 
find that an element, which in the free state we regard as 
a metal, in combination acts like a non-metal. 

Of the seventy-five or more elements now generally 
recognized, about twenty are regarded as non-metals, the 
rest being metals. Of the latter more than one-half are 
rare, and, in several cases, so rare that some doubt exists 
as to whether they are actually elements, or mixtures of 
elements. 



178 LECTURE NOTES 

221. History and Occurrence. The ancients knew but 
six of the metals, viz. : gold, silver, copper, iron, tin and 
lead. They knew the alloy of copper and zinc, called 
brass, but did not know zinc as a separate metal. They 
also knew the non-metals carbon and sulfur. These are 
all mentioned by the writers of the Old Testament, as 
well as by the early Greek writers. Mercury was first 
mentioned by Theophrastus about B. C. 300, and zinc, 
arsenic, antimony and bismuth were known to the early 
alchemists. Most of the other metals have been discov- 
ered since the middle of the eighteenth century. 

All of the metals are found in the earth's crust, although 
only a few of them are found free. They are usually found 
combined with oxygen or sulfur alone, or with oxygen and 
some other of the non-metallic elements, and form the 
natural compounds which we call minerals. The greater 
portion of the minerals are found mixed together more or 
less closely, and form the rocks which constitute the 
greater portion of the earth's crust. The minerals from 
which the metals are obtained are called ores. While a 
large number of minerals may be used for obtaining the 
metals, comparatively few are commonly used for this 
purpose. The most common ores are the oxids, sulfids 
and carbonates, and, in a few cases, the chlorids, sulfates, 
or other salts. 

222. Preparation. The process of obtaining a metal 
from its ores is often quite complicated, and the study of 
the different processes employed forms that branch of nat- 
ural science called metallurgy. 

In general the ores sought are the oxids. If other ores 
are used they are, if possible, first converted into oxids. 
This is done by heating the ore in contact with air in a 



LECTURE NOTES 179 

properly constructed furnace, the process being called 
roasting, thus: 

CuS + 3 O = CuO + S0 2 . 
ZnCO s =ZnO + C0 2 . 

The oxids are usually reduced to the metal by means of 
carbon, the operation frequently requiring a very high 
temperature. The following equation shows the action 
upon ferric oxid : 

Fe 2 O s -f- 3 C = 2 Fe + 3 CO, 

The carbon monoxid thus formed also helps in the re- 
duction, thus: 

Fe 2 3 + 3 CO = 2 Fe + 3 CO,. 

The chlorids are sometimes reduced by iron or copper 
and in the case of some of the rarer elements by means of 
hydrogen or metallic sodium, thus: 

MgCl 2 + 2 Na = Mg + 2 NaCl. 

At the present time a number of metals are obtained 
by electrolytic processes, and these bid fair to replace 
entirely many of the older and more expensive methods. 
The particular methods employed will be described under 
the individual metals. 

223. Physical properties of the Metals. The physical 
properties of the metals are many of them such as can be 
determined by simple observation, while others have ref- 
erence to their relations to mechanical force, and still 
others to the effects of heat and electricity. 

At the ordinary temperature all metals are solid, except 
mercury, which is liquid. 

The metals all have what is called metallic luster, that 
is, they reflect light from their surface. 



180 LECTURE NOTES 

The metals vary much in hardness. Nickel and steel 
are very hard, lead may be cut with a knife, sodium and 
potassium are quite soft, and mercury is a liquid. 

The color of the metals is generally light, and varies 
from silver white to lead gray. A few are more highly 
colored, copper being red, while gold, calcium, and stron- 
tium are yellow. In the form of powder most of the metals 
are black. 

The metals are all opaque as ordinarily seen. A few, 
in the form of very thin foil, transmit some light and so 
are semi-transparent. Thus gold foil transmits green 
light and silver foil blue light. 

Most of the metals, when pure, show a crystalline struc- 
ture, the crystals usually being regular in form. 

The specific gravity of the metals varies greatly. The 
specific gravity of lithium is only 0.59, while that of os- 
mium is 22.47. The most important ones have a specific- 
gravity between 7 and 14, those above 5 being called 
heavy metals. The heavy metals usually show rather 
weak metallic properties when combined with other 
elements. 

The following properties are dependent upon the rela- 
tions of the metals to mechanical, and other forces, and 
are quite important with reference to the commercial 
value of the metals. 

Metals are generally malleable and so can be hammered 
into thin sheets. Gold is the most malleable metal, and 
gold-leaf is usually about 0.0001 m. m. in thickness, the 
extreme limit being about 0.00008 m. m. Antimony and 
bismuth are quite brittle, and can be powdered in a 
mortar. 



LECTURE NOTES 181 

Most metals are ductile, that is, can be drawn out into 
a wire. Gold is the most ductile metal, and two miles of 
the finest gold wire weigh only about one gram. 

Most of the metals are tenacious, that is, are capable 
of resisting strain. Some of them, especially nickel, and 
iron in the form of steel, possess this property in a very 
marked degree. 

The metals are all conductors of heat and electricity. 
Silver and copper are the best conductors. Copper is 
universally employed for this purpose on account of its 
relative cheapness. 

The melting point of the metals varies greatly. Mer- 
cury melts at — 39°, platinum at about 2000°, a few rare 
metals have a still higher melting point, and osmium is 
practically infusible. 

224. Specific Heat. If we subject the same weights of 
different metals to the same source of heat, different 
lengths of time are required in which to raise them to the 
same temperature. This varying heat capacity of metals 
is compared with the heat capacity of water; and the 
quantity of heat necessary to raise the temperature of a 
certain weight of any substance one degree, as compared 
with the quantity of heat necessary to raise the tempera- 
ture of the same weight of water one degree, is called the 
specific heat of the substance. 

If we mix one kilo of water at 0°, and 1 kilo of mercury 
at 100°, the mixture will have a temperature of 3.1°. 
This means that the total loss of heat by the mercury, 
just equals the total gain of heat by the water, and that 
the numbers representing the temperatures lost and gain- 
ed, are inversely proportional to their heat capacities. The 



182 



LECTURE NOTES 



specific heat of water being taken as 1, we find the specific 
heat of mercury by dividing 3.1 by 96.9, which gives 
0.0319, as the specific heat of mercury. The following 
table gives the specific heat of a few of the more impor- 
tant metals: 



Aluminum, 


0.2140 


Magnesium, 


0.2500 


Antimony, 


- 0.0523 


Manganese, - 


- 0.1220 


Arsenic, - 


0.0822 


Mercury, 


0.0319 


Bismuth, 


- 0.0305 


Nickel, 


- 0.1080 


Cadmium, 


0.0567 


Platinum, 


0.0324 


Copper, 


- 0.0952 


Silver, 


- 0.0570 


Iron, 


0.1140 


Tin, 


0.0548 


Lead, - 


- 0.0315 


Zinc, - 


- 0.0955 



225. Atomic Heat. If we compare the specific heats of 
the elements with their atomic weights, we shall see that 
the numbers representing them are inversely proportional 
to each other. If we multiply the specific heat of each 
element by its atomic weight, we shall obtain a constant 
number, or, at least, a series of numbers that are very 
nearly the same. The average of these numbers, which 
is 6.4, is known as the atomic heat. This remarkable 
fact, that all elements possess the same, or approximately 
the same, atomic heat, is known from its discoverers as 
the law of Dulong and Petit. 

Since the specific heat may be determined by direct ex- 
periment, it becomes a means of determining the atomic 
weight of an element. For, from the above statement, 
the atomic weight of an element is equal to 6.4 divided by 
its specific heat. 

226. Variations in Atomic Heat. It was early noticed 
that, in the case of certain elements, particularly some of 



LECTURE NOTES 183 

the non-metals, the atomic heat obtained at ordinary- 
temperatures, was much lower than 6.4. The following 
table shows the atomic heats of these elements: 



Diamond, 


1.8 


Glucinum, 


3.7 


Gas carbon. 


- 2.2 


Silicon, 


- 4.9 


Graphite, 


2.4 


Sulfur, - 


5.2 


Boron, 


- 2.6 


Phosphorus, 


- 5.4 



Investigation showed that the specific heat of the dia- 
mond increased rapidly with the temperature, and that 
at 1000° the atomic heat became 5.5. The same is true 
for the other exceptional elements. The elements which 
show the greatest variation all have a more or less com- 
plex molecular structure, and are exceptional in other 
ways. Thus, all forms of carbon as well as boron and silicon 
are infusible. These variable elements all have atomic 
weights which are less than 35, and while many elements 
of low atomic weight conform to the law, the exceptional 
elements are all found here. 

The reason for the variation is not definitely known. 
Heat is a very complex force. It not only causes rise of 
temperature in bodies, but does work in them in various 
ways. The amount of heat required for this work is 
variable, and this is therefore believed to be the principal 
cause of the variation. 

227. Chemical Properties of the Metals. The relations 
of the different elements to oxygeu form the best means of 
distinguishing them as metals and non-metals; for, as we 
have already noticed (33), the non-metals when combined 
with hydrogen and oxygen form acids, while the metals 
under similar conditions form bases. Only a few of the 
bases are soluble in water, and so only a few of them give 



184 LECTURE NOTES 

the basic or alkaline reaction with vegetable colors, such 
as litmus, but all bases show their character by combin- 
ing with acids to form salts. 

It not infrequently happens that a metal in combination 
with other elements, is so influenced by them as to act 
like a non-metal. This usually happens when the metal 
is combined with an excess of oxygen, and especially so 
when in the presence of some very strongly basic metal, 
such as the so-called alkali-metals. 

As a rule, those metals which have a low specific gravity 
form the strongest bases, and. conversely, the heavy 
metals form weak bases. 

228. Alloys. Alloys are mixtures of the metals, fused 
together so as to form an apparently homogeneous mass. 
They may be regarded as solutions of solids in solids. Not 
all metals will form alloys when fused together, although 
the most of them will do so. From the fact that certain 
alloys have a distinctly crystalline structure, it is believed 
by some that there is a weak chemical union between the 
metals forming them. The careful examination of a large 
number of alloys seems to indicate that this is not true, 
although in a few cases undoubted evidence of chemical 
combination has been observed. 

When an alloy, which in the liquid condition is perfectly 
homogeneous, is cooled, it very often happens that the 
solidified mass is no longer homogeneous, but that differ- 
ent proportions of the metals are found in different parts 
of the solid. This is known as the liquation of alloys. 

The melting point of an alloy is usually lower than that 
of any of the metals contained in it. Thus, common 
solder melts at atemperature lower than either the lead 



LECTURE NOTES 185 

or tin of which it is composed. A remarkable example 
of this is the alloy known as ; 'Wood's metal." This is 
composed of four parts of bismuth, two of lead, one of tin, 
and one of cadmium, and melts at 65°, although the tin, 
which has the lowest melting point of any of the metals, 
melts at 232°. 

Many of the other physical properties are changed when 
the metals are alloyed. Brass, which is composed of 
copper and zinc, is more tenacious than either metal alone. 
Both gold and silver are too soft to be used for coinage, 
but when alloyed with copper, also a soft metal, they 
each have the necessary hardness. 

The solubility of the metals in acids is often changed 
when they are alloyed. Platinum is insoluble in nitric 
acid, but alloyed with silver it dissolves easily in that 
acid. Silver dissolves easily in nitric acid, but when 
alloyed with 33 per cent, of gold, it is entirely insoluble 
in the acid. 

In some alloys the electric conductivity is a mean of that 
of the metals of which they are composed, while in others, 
especially those of gold, silver, copper and zinc, the con- 
ductivity is less than a mean of the component metals. 

229. Some Important Alloys. Although all the metals 
can form alloys, only a few 7 of them are used to any extent 
for this purpose. Of these the most important are gold, 
silver, copper, nickel, tin, zinc, lead. iron, antimony, bis- 
muth, aluminum and cadmium. 

The first four of these are universally employed for 
coinage. G-old and silver coins are nearly always alloyed 
with 10 per cent, of copper, although in England the gold 
and silver coins contain 8.34, and 7.5 per cent, of copper 



186 LECTURE NOTES 

respectively. Nickel coins contain 75 per cent, of copper. 
The smaller coins are usually made of bronze. 

Brass is an alloy of two parts of copper and one of zinc. 
Copper is so soft as to be useless for many purposes, but 
brass is much harder and can be turned in a lathe. 

Gun-metal is made of nine parts of copper and one of 
tin, and is very hard and tenacious. 

Bell-metal is made of four parts of copper and one of 
tin and is much harder. Other metals are sometimes 
added to improve^the tone of the bell. 

Speculum is made of two parts of copper and one of tin. 
It is almost white, is very hard, and takes a high polish. 
It is used for making telescopic mirrors. 

Bronze is made of varying quantities of copper and tin, 
to which is added a little zinc and sometimes lead. It is 
hard and tough and can be worked in a lathe. 

Aluminum brotize is made of copper with 10 per cent, 
of aluminum, and in color much resembles standard gold. 

Type-metal is made of four parts of lead and one of 
antimony. It is hard, easily fusible, and expands slightly 
when it solidifies, making thus a perfect cast. 

Britannia, pewter, and the white or anti-friction metals 
are made of tin, antimony and small quantities of copper. 
One of these, called Babbit's metal, contains about 40 per 
cent, of lead in place of part of the tin. 

230. Amalgams. Amalgams are alloys of the metals 
with mercury. Most of the metals form amalgams, which 
can usually be obtained by direct union. Heat is some- 
times evolved during the process, especially in the case 
of the strongly basic metals such as sodium and potassium. 



LECTURE NOTES 187 

In a few cases an amalgam can be formed by adding mer- 
cury to a solution of a metallic salt, or by adding a metal 
to a solution of mercuric nitrate. 

If the mercury is in large excess, the amalgam may be 
liquid, but most of them are solids and they are often 
crystalline. 

When an amalgam is heated, the mercury is partly or 
entirely driven off, a fact which is utilized in the amalga- 
mation process for extracting gold from its ores. 

Many amalgams find use in the arts. Tin amalgam is 
used for coating the backs of mirrors. Amalgams of tin 
and silver are used for filling teeth. Sodium amalgam is 
a valuable reducing agent, and is much used in organic 
chemistry. 

231. The Oxids and Hydroxids of the Metals. All of 

the metals form compounds with oxygen, called oxids. 
and most of them form compounds with oxygen and hy- 
drogen, called hydroxids. A few of the metals have a 
very strong affinity for oxygen, so that they oxidize easily 
in moist air, and decompose water at the ordinary tem- 
perature. The oxids of these metals combine directly 
with water to form the hydroxids, which are very strong- 
bases, and, being soluble in water, are called alkalies 
(38). These strong bases are not usually decomposed by 
ignition. 

Some of the heavy metals decompose water at a high 
temperature forming oxids. Most of these oxidize slowly 
in moist air at the ordinary temperature. They do not 
usually form hydroxids directly, but the hydroxids can 
be formed by precipitating a solution of one of their 
salts by one of the soluble bases, thus : 

MgS0 4 -f 2 KOH = Mg(OH) 2 + K 2 SO + . 



188 LECTURE NOTES 

These hydroxids are all decomposed by heat into the oxid 
and water, thus: 

Mg(OH) 2 = MgO + H 2 0. 

A few of the heavy metals such as gold, silver, plati- 
num, etc., cannot be made to combine with oxygen 
directly, and for this reason are called the noble metals. 
Their oxids can be obtained indirectly by precipitating a 
solution of one of their salts by a soluble base, thus: 

2 AgNO s + 2 NaOH == 2 NaN0 8 + Agfi -f H,0. 

The oxids "of the noble metals are all decomposed by 
heat into the metal and oxygen. 

All these normal oxids and hydroxids exhibit their basic 
character by combining with acids to form salts, and 
water, thus: 

Mg(OH) 2 + H 2 S0 4 = MgS0 4 + 2 H 2 0. 

A few of the less basic oxids and hydroxids combine 
with the strong soluble bases to form salts, in which they 
exhibit a weak acid character, thus: 

Al(OH) 8 + 3 NaOH = Al (ONa) 3 + 3 H 2 0. 

232. The Higher Oxids. Most of the metals are capa- 
ble of a higher degree of oxidation than that which is 
found in the normal oxids, which are formed by the 
methods just given. These higher oxids are called 
per oxids. 

There are two classes of peroxids: first, those of the 
more basic metals, which, when acted upon by dilute acids, 
yield hydrogen dioxid, and do not form salts with the 
soluble bases ; second, those which do not yield hydrogen 
dioxid and generally do form salts with the soluble bases. 
All the peroxids, when strongly ignited, give off oxygen. 



LECTURE NOTES 189 

They also give off oxygen when warmed with concentrated 
sulfuric acid, and when heated with concentrated hydro- 
chloric acid give off chlorin. Some of these higher oxids 
exhibit a fairly strong acid character. 

233. The Halogen Compounds of the Metals. All of the 

metals form compounds with the halogen elements, the 
most of which are very stable. They will usually combine 
with the halogen directly, but this direct method, while 
of theoretical importance, is rarely used. 

The halogen compounds are usually made by dissolving 
the metals themselves, their oxids, hydroxids, or carbon- 
ates, in the corresponding halogen acid, thus: 

Zn -f 2 HC1 = ZnCl 2 -f H 2 
ZnO -f 2 HC1 = ZnCl 2 + H 2 0. 
ZnCO a -f 2 HC1 = ZnCl 2 + H 2 -f C0 2 . 

It is by reason of this action that the halogen com- 
pounds of the metals, although they are binary compounds, 
are regarded as salts of the corresponding halogen acids. 

The insoluble halogen compounds are made by acting 
upon a solution of a soluble salt of the metal, with a solu- 
tion of a soluble salt of the corresponding halogen ele- 
ment, thus: 

AgN0 3 + KBr = AgBr -f KNO g 
Pb ( N ° 3 ) 2 +2KI = Pbl 2 + 2 KN0 3 . 

A few of the halogen compounds are made by special 
methods. 

234. The Other Metallic Salts. When a metal is acted 
upon by any acid, the hydrogen of the acid is replaced by 
the metal forming a salt, and setting free the hydrogen. 
[The nascent hydrogen thus set free, may, and often 



190 LECTURE NOTES 

does, reduce some of the acid, setting free some other 
gas. (See 101 and 131.) ] 

Since all metals are not acted upon by all acids, this 
method of forming salts can only be used in a limited 
number of cases. As already noted (231), the oxids or 
hydroxids are acted upon by the acids forming salts, and 
water, the metal in the base replacing the hydrogen in 
the acid, thus: 

Mg(OH) 2 + H 2 S0 4 = MgS0 4 + 2 H 2 0. 

Instead of the oxids or hydroxids, some weaker salt, 
such as the carbonate, may be decomposed by an acid 
forming the corresponding salt of the acid, thus: 

CaC0 3 + 2 HNO s = Ca(N0 3 ) 2 + H 2 -f C0 2 . 

The insoluble salts may be formed by acting upon a 
solution of a soluble salt of the metal, by a solution of a 
soluble salt of the acid. The following equations will 
illustrate this: 

BaCl 2 + Na 2 S0 4 = BaSC> 4 + 2 NaCl 
AgNO s + NaCl == AgCl + NaN0 3 . 

235. The Natural Groups of the Elements. In our study 
of the non-metals we have seen that certain of them seem 
to fall into groups, according to their chemical properties, 
and that, in many cases, members of the same group bear 
a striking resemblance to each other. Marked examples 
of this are seen in the so-called halogen group, and the 
sulfur group. 

If we take the three important members of the halogen 
group, viz. : chlorin, bromin and iodin, we shall see that 
there also exists some relation between their atomic 
weights and their properties, for the atomic weight of 



LECTURE NOTES 191 

bromin is very nearly a mean of that of chlorin and iodin. 
This is also true of the chemical and physical properties 
of bromin. The same relation exists in the three impor- 
tant members of the sulfur group, viz. : sulfur, selenium 
and tellurium. As we study the metals we shall see that 
they form similar groups. Thus the metals lithium, so- 
dium and potassium resemble each other very closely, as 
do also calcium, strontium and barium. 

A careful study of these and other similar groups led 
to the important generalization, that there exists some 
relation between the chemical and physical properties of 
all the elements and their atomic weights. This resulted 
in the so-called periodic classification (34), to which ref- 
erence has already been made. 

236. The Periodic System. If we arrange all the ele- 
ments in the order of their increasing atomic weights, we 
shall find that those elements which have been recognized 
as belonging to the same natural group, recur in regular 
periods. 

The first element is hydrogen, which is an element be- 
longing to no natural t ; group. ,_] This element possesses 
properties which are unique, and, together with helium, 
forms the first series, and what we may call the introduc- 
tory period. It very possibly may develop that among the 
very recently discovered elements, some will be found to 
belong to this series, but at the present time too little is 
known concerning them to warrant any positive state- 
ment. 

[In order to make the relationship of the elements more 
apparent, we will at first omit in the arrangement, those 
gases recently discovered in the atmosphere (120).] 



192 



LECTURE NOTES 



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LECTURE NOTES 193 

. After helium we find the next seven elements are lith- 
ium, glucinum. boron, carbon, nitrogen, oxygen and 
fluorin. This is the first or typical period. The next 
seven elements, which form the second period, are sodium, 
magnesium, aluminum, silicon, phosphorus, sulfur and 
chlorin. If these two periods are arranged in two hori- 
zontal rows, the elements falling in vertical columns will 
be found to be very similar in their properties, thus : 

Li Gl B C N O F 

Na Mg Al Si P S CI 

The next period, beginning with potassium, has for its 
seventh element manganese, an element which forms an 
oxid analagous to that of chlorin. The three elements 
following, viz. : iron, nickel and cobalt, are regarded as 
belonging to an eighth group, and as connecting manga- 
nese with the next element copper, which forms an oxid 
corresponding to those of sodium and potassium. Con- 
sidering copper, therefore, as the first of the fifth series, 
the seventh element is bromin, and this third period is 
then complete. This third period is therefore double the 
size of the first two periods, and has the connecting eighth 
group of three elements. The next, or fourth period is 
similar, consisting of two series and three connecting 
elements belonging to the eighth group. The higher 
periods have points of resemblance, but are far from com- 
plete. They include the so-called rare earths, about 
which very little is certainly known. [See table on 
opposite page.] 

Some years ago, Thomsen pointed out that in the sudden 
change from the very strongly negative elements, such 
as fluorin. chlorin, bromin, and iodin, to the very strongly 
positive elements which follow each of them, viz.: sodium, 



194 LECTURE NOTES 

potassium, rubidium, and caesium, there should be in each 
case a connecting element, which should be electrically- 
indifferent, and whose valence should be zero, that is, it 
should form no compounds. He also predicted the atomic 
weight of each of these elements. The recent discovery of 
the four gases in the atmosphere, viz. : neon, argon, 
krypton and xenon, and the determination of their atomic 
weights by Ramsey (120), seems to show that they are 
the predicted elements, belonging to the eighth group. 

In the table given on page 102, the blaDk spaces 
in the higher periods probably represent elements as yet 
unknown. Three blank spaces were found in the third 
period when this arrangement was first proposed. The 
properties of each unknown element were predicted from 
their relations to the surrounding elements, and when 
scandium, gallium and germanium were discovered, they 
were found to agree exactly with the predicted elements. 

237. The Relations of the Chemical and Physical Prop- 
erties to the Periodic System. If we study the elements 
in the first group, we shall find that they include the 
most strongly electro-positive elements known, while the 
seventh group contains the halogen elements, which in- 
clude the most strongly electro-negative elements known. 
The elements in the other groups, except the eighth, 
occupy an intermediate position. 

We also find that as the atomic weight increases, the 
periods show an increase in the basic character of their 
members. Thus the first short period contains five non- 
metals, viz. : boron, carbon, nitrogen, oxygen, and fluorin; 
the second period contains four, viz. : silicon, phosphor- 
us, sulfur, and chlorin; the first long period three, viz.: 
arsenic, selenium, and bromin; the second long period but 



LECTURE NOTES 195 

two, viz.: tellurium and iodin; and the higher periods 
none as yet known. 

If we examine the valence of the elements which belong 
to the two short periods, we shall see that when the ele- 
ments are combined with hydrogen, or with elements 
with which they exhibit their hydrogen valence, this 
valence is one in the first group, increases to four in the 
fourth group, and decreases to one in the seventh group. 
When the elements combine with oxygen, and exhibit their 
highest valence, the valence is one in the first group, and 
increases regularly to seven in the seventh group. In 
the first four groups, therefore, there is no variation in 
the valence. 

The elements which occur in the middle of the long 
periods, and which belong to the eighth group, do not, as 
a whole, show any very characteristic valence. From 
analogy we should expect them to have a valence of eight 
in their highest oxygen compounds. Only two of the 
elements belonging to this group, show this very high 
valence. These are ruthenium and osmium, which form 
the compounds Ru0 4 and Os0 4 , respectively. The others 
form various oxids, and do not form compounds with 
hydrogen. 

The elements of this group form a series of characteris- 
tic compounds in which they form very complex acid 
radicals with cyanogen, and another series in which they 
form complex basic radicals with ammonia. They also 
have other rather exceptional properties. 

In the higher periods there are some apparent discrep- 
ancies in valence. We must, however, remember, not 
only that the higher periods are as yet very incomplete, 
but also that the valence of an element depends upon the 



196 LECTURE NOTES 

nature of the combining elements. With this in mind, 
we shall see that in the higher periods the elements also 
show the same periodicity in their valence. 

Specific gravity is highest in the middle of the periods 
and decreases to the ends. Other important properties, 
such as the melting point, boiling point, malleability, and 
conductivity, also exhibit periodicity. 

Since the atomic weight of an element determines its 
position in the periodic system, the above relations show, 
that the properties of the elements are periodic functions 
of their atomic weights. 

238. The Alkali-Metals. The two metals, sodium and 
potassium, were first made from compounds obtained by 
leaching ashes, which compounds were called alkalies. 
[From the arabic al qali, the ash.] Other metals were 
discovered having properties similar to these two, and 
now there are five such known. They are lithium, sodium, 
potassium, rubidium and caesium. 

All the alkali-metals oxidize easily in moist air, decom- 
pose water at the ordinary temperature, liberating hydro- 
gen, and form strongly basic hydroxids which are called 
alkalies, or caustic alkalies. Few of their compounds are 
decomposed by heat. All of their compounds are soluble 
in water to some extent, and most of them quite easily. 
They all belong in Group I, of the periodic system, and 
are therefore all monovalent. They are the most strongly 
positive elements known. 

[The compound radical ammonium, NH 4 , forms com- 
pounds exactly analogous to those of this group, and so is 
sometimes called the volatile alkali, since all of its com- 
pounds volatilize at a comparatively low temperature. 
The other members of the group are called fixed alkalies.] 



LECTURE NOTES 197 

In taking up the study of the metals and their com- 
pounds, only those which are common enough to be of 
some commercial importance, will be described at length. 

Lithium. Symbol Li. Atomic Weight 7. 

239. Occurrence, Preparation, and Properties. Lithium 
is found in nature, quite widely distributed, but in small, 
quantities. It never occurs free, but its compounds are 
found in a number of minerals, the most important of 
which are lepidolite, amblygonite, and spodumene. It also 
occurs in a number of mineral springs, and in many plants, 
such as coffee, sugar-cane, tobacco, and beets. 

Lithium can be prepared by electrolysis of the fused 
chlorid. and by other processes similar to those by which 
sodium and potassium are prepared. [See 242 and 257.] 

Lithium is a silver-white metal, which tarnishes slowly 
in the air, and decomposes water at the ordinary temper- 
ature. It is the lightest metal known, its specific gravity 
being only 0.59. It melts at 180°, and, at a high tem- 
perature, burns with a bright red flame. 

240. The Compounds of Lithium. The compounds of 
lithium are quite numerous, but not of great importance. 
They must resemble the corresponding compounds of 
sodium. 

Lithium carbonate, Li 2 C0 3 , is formed when a solution 
of any of the other alkaline carbonates is added to a solu- 
tion of lithium chlorid or nitrate. It is only slightly 
soluble in water. It is used in medicine as a remedy for 
gout and rheumatism, as it acts as a solvent for uric acid. 

Lithium phosphate, LiJPO , is almost entirely insoluble 
in water. The other salts are formed by the general 
methods, and are all easily soluble in water. 



198 LECTURE NOTES 

All lithium eompoimds color the non-luminous gas flame 
a beautiful red color. 

Sodium. Symbol, Na. (Natrium). Atomic Weight, 23. 

241. History and Occurrence. Compounds of sodium 
have been known since the very earliest times, but the 
metal itself was first obtained by Davy, in 1807. 

Sodium is the eighth most widely distributed element 
known. It constitutes about 2.3 per cent, of the earth's 
crust. It does not occur free in nature, but its compounds 
are very numerous and very widely distributed. Its most 
important compound is the chlorid, NaCI, called common 
salt, which is found in sea- water, in salt lakes and springs, 
and as rock-salt in the numerous salt deposits of the world. 
Sodium nitrate, NaNO g , called also Chili saltpeter, is 
found in large deposits on the west coast of South 
America. 

Sodium is found in considerable quantities in the rocks, 
especially in cryolite and some of the feldspars, and in 
small quantities in a very large number of minerals. It 
is found in many animal and vegetable substances, espe- 
cially in those occurring in or about the sea. 

242. Preparation. Sodium was first prepared by Davy 
by electrolysis of sodium hydroxid. Some years later a 
process was devised by Brunner by which the metal was 
obtained by strongly igniting a mixture of sodium car- 
bonate and carbon, thus: 

Na 2 CO s + 2 C = 2 Na + 3 CO. 

This process was employed until 1886, when it was 
improved by Castner, who replaced the sodium carbonate 
with sodium hydroxid. The action is as follows: 

6 NaOH + 2C^2 Na 2 C0 3 -f- 3 H 3 + 2 Na. 



LECTURE NOTES 199 

In 1890, Castner introduced the present electrolytic 
method, which is only an improvement of Davy's original 
method. This was made practical by the improvement in 
electrical machinery, and has resulted in a very great re- 
duction in the price of sodium. The apparatus is quite 
simple, and is so arranged that the process is continu- 
ous. The temperature must be kept fairly constant at 
just above the melting point of sodium hydroxid, viz. : 
310°. 

All the sodium of commerce is now made by this process. 

243. Properties and Uses. Sodium is a silver-white 
metal, and has a bright metallic luster. It oxidizes quick- 
ly in moist air, and so is usually kept under some hydro- 
carbon oil, although it may be preserved for a long time 
in perfectly dry air. If it is melted, and allowed to oool 
in an atmosphere of hydrogen, it can be obtained in crys- 
tals. It decomposes water in the cold, liberating hydro- 
gen. If the water is hot, or if the metal is confined to 
one place, the heat of the chemical action will ignite the 
hydrogen, which then burns with a characteristic yellow 
flame. 

Sodium has a specific gravity of 0.97. It melts at 
95.6° and boils at 742°, giving off a vapor which is violet 
by transmitted light. 

Sodium is used quite extensively in the manufacture of 
certain elements and compounds, such as silicon, boron, 
magnesium, certain cyanids, and sodium peroxid. It is 
also much used in the laboratory as a reducing agent, 
both as metal, and as sodium amalgam. 

COMPOUNDS OF SODIUM. 

244. Sodium Hydroxid, NaOH. This important com- 
pound, which is also known as caustic soda, may be ob- 



200 LECTURE NOTES 

tained by adding metallic sodium, in small pieces, to a 
little water in a silver dish, and evaporating to dryness. 

Until quite recently it was always made commercially 
by boiling a solution of sodium carbonate with calcium 
hydroxid (slaked lime), thus: 

Na 2 CO s + Ca(OH) 2 = CaC0 3 + 2 NaOH. 

An electrolytic process has recently been devised by 
Castner, which is likely to replace all other processes of 
manufacture, because of its cheapness, and the high 
degree of purity in the product. The process consists in 
decomposing a solution of sodium chlorid by an electric 
current. The sodium, which is formed at the negative 
pole, decomposes the water, forming the hydroxid and 
hydrogen, while chlorin is liberated at the positive pole. 
In order to avoid the action of the chlorin upon the sodium 
hydroxid (which would form sodium chlorate and sodium 
hypochlorite), the decomposing cell is divided into three 
compartments, which are separated by a mercury dia- 
phragm. The outer compartments contain the salt 
solution, and the positive pole, the inner one water, and 
the negative pole. The sodium first forms an amalgam 
with the mercury, which is decomposed by the current at 
the negative pole, and then the sodium combines with the 
water forming the hydroxid. 

Sodium hydroxid is a white crystalline solid. It melts 
at 310°, and volatilizes at a high temperature. It is very 
deliquescent, and, when moist, absorbs carbon dioxid from 
the air, forming sodium carbonate. The fused substance, 
or even a strong solution of it, attacks most metals to 
some extent, silver being least affected by it. 

Sodium hydroxid is extensively used in the laboratory 



LECTURE NOTES 201 

for neutralizing acids, and for precipitating the hydroxids 
and oxids of the heavy metals. 

245. Sodium Dioxid, Na 2 2 . This compound, which is 
also called sodium peroxid, is obtained by burning sodium 
in oxygen. 

Sodium dioxid is a yellowish- white powder, which decom- 
poses when dissolved in water, forming sodium hydroxid 
and oxygen. If the water is kept cold, hydrogen dioxid 
is formed. It is a very powerful oxidizing agent, so 
much so, that in contact with easily combustible sub- 
stances, such as wood, paper, etc.. the action is often so 
energetic as to set them on fire. 

It is used commercially for bleaching straw, under the 
name of soda-bleach. 

246. Sodium Chlor id, NaCl. This very important com- 
pound, which is known commercially as salt, occurs in 
immense quantities in sea- water, and in salt lakes and 
springs As rock-salt it occurs in vast deposits in many 
parts of the earth, especially at Stassfurt. in Germany, 
and in central New York. 

It may be made by the direct union of sodium and 
chlorin, or by acting upon metallic sodium, its hydroxid, 
or carbonate, with hydrochloric acid. For commercial 
and domestic purposes, it is always obtained either by 
evaporation of salt water, by mining the rock-salt, or by 
means of artificial salt-wells. The latter are made by 
boring into the salt deposit and allowing water to pass 
into the opening. The water soon becomes saturated 
with the salt and is pumped out. and the salt obtained 
by evaporation. 

Sodium chlorid is a white crystalline solid, the crystals 



202 LECTURE NOTES 

being regular cubes. It has an agreeable saline taste, 
and is an important article of food. It melts at 815°, and 
volatilizes at a temperature a little above this. The pure 
compound does not attract moisture, but as the common 
salt contains traces of magnesium salts, it is always 
somewhat deliquescent. 

Sodium chlorid is used in a large number of commercial 
processes, such as the manufacture of sodium carbonate, 
soap, paper, and glass, as well as in meat-packing and in 
the dairy business. It is also the principal source of 
supply for both sodium and chlorin, as well as for the 
compounds of both. 

247. Sodium Sulfate, Na 2 S0 4 . This substance in its 
crystalline form is also called Glauber's salt. It is quite 
widely distributed in nature, being found in numerous 
deposits, and in springs. It is produced in large quanti- 
ties in the manufacture of soda, being there known by the 
commercial name of salt-cake. 

Sodium sulfate forms large colorless prismatic crystals, 
having ten molecules of water of crystallization, which is 
expressed by the formula, Na- 2 S0 4 , 10 H 2 0. It is very 
soluble in water, the temperature of its greatest solubility 
being 34°, at which temperature the crystals lose their 
water of crystallization, forming the anhydrous salt, 
Na 2 S0 4 . The latter salt melts at 863°. The crystals also 
part with their water of crystallization when exposed to 
the air, and are, therefore, efflorescent. 

Sodium sulfate exhibits in a very striking manner the 
phenomena of supersaturated solutions (22); for if a solu- 
tion saturated at 34° is allowed to cool, no separation of 
crystals occurs until some solid body, preferably a crys- 



LECTURE NOTES 203 

tal of the substance, is introduced, when the whole 
suddenly crystallizes, with considerable evolution of heat. 
[See Latent Heat, 20.] 

248. Sodium Nitrate, NaN0 3 . This compound, which 
is known as Chili saltpeter, is found in immense deposits 
in parts of Chili, Peru, and Bolivia, where it is also known 
as caliche. These deposits contain from 25 to 65 per cent, 
of sodium nitrate, the rest being sodium chlorid, sodium 
sulfate, sodium iodate, with some potassium nitrate and 
calcium sulfate. From this the pure sodium nitrate is 
obtained by solution and crystallization. 

Sodium nitrate is a white crystalline solid. It melts 
at 318°, is somewhat deliquescent, and easily soluble in 
water. It is used for the manufacture of nitric and sul- 
furic acids, and as a fertilizer, since it readily gives its 
nitrogen to plants growing in a soil which contains it. 

249. The Phosphates of Sodium. Sodium forms three 
phosphates according as one, two, or three hydrogen 
atoms in phosphoric acid are replaced by sodium, viz. : 
primary, or monosodium phosphate, Na 2 HP0 4 ; secondary, 
or disodium phosphate, NaH 2 P0 4 ; and tertiary, or tri- 
sodium phosphate, Na 3 P0 4 . 

The secondary sodium phosphate, Na 2 HP0 4 , is the 
most common and important of the series, aud is always 
referred to by the simple name sodium phosphate. 

It is formed by neutralizing phosphoric acid by means 
of sodium carbonate or hydroxid. It is quite soluble in 
water, from which solution it crystallizes, at the ordinary 
temperature, in large transparent crystals, which contain 
twelve molecules of water of crystallization, and effloresce 
rapidly in the air. When heated, the salt loses water, 



204 LECTURE NOTES 

melts at about 300°, and forms sodium pyrophosphate, 
Na 2 P a O . It is much used in medicine. 

250. Sodium Carbonate, or Soda, Na 2 C0 3 . This impor- 
tant compound is known commercially as carbonate of 
soda, or soda-ash. It occurs in large quantities in the 
so-called soda-lakes found in many parts of the earth, 
especially in Egypt, Central Africa, South America, and 
California. In Egypt this natural soda is called trona, 
and in South America, urao. It is manufactured on an 
enormous scale by the processes described below. 

Sodium carbonate crystallizes at the ordinary tempera- 
ture, with ten molecules of water of crystallization, 
Na 2 C0 3 , 10 H 2 C At different temperatures, the amount 
of water in the crystals varies. The crystals are efflores- 
cent. The anhydrous salt melts at 850° and volatilizes 
at a very high temperature. It has a strong alkaline 
reaction, although it is a normal salt. Nearly all acids 
decompose it, liberating carbon dioxid. 

251. The Leblanc Process. Previous to 1789, all the 
sodium carbonate was obtained by the leaching of kelp, 
or varec (83). During the French Revolution, the supply 
being cut off, the French government offered a prize for 
the best method of manufacturing this important sub- 
stance, which was won by a chemist named Leblanc. 
Practically all the methods known employ sodium chloric! 
as the source of the sodium. 

In the Leblanc process the sodium chlorid is first acted 
upon by sulfuric acid, forming sodium sulfate, thus: 

H a S0 4 + 2 NaCl = Na 2 S0 4 + 2 HC1. 

This part of the work is known as the salt-cake process. 

The sodium sulfate is then mixed with carbon (powdered 



LECTURE NOTES 205 

coal), and calcium carbonate (limestone), and strongly 
ignited in a reverberatory furnace. This part of the work 
is known as the black-ash process. 

The action is two-fold. In the first action the carbon 
reduces the sodium sulfate, thus: 

Na 2 S0 4 + 2 C = Na 2 S + 2 C0 2 . 

The sodium sullid now reacts upon the limestone, thus: 

Na 2 S + CaC0 3 = CaS + Na 2 C0 3 . 

The black-ash is then subjected to a washing process, 
called lixiviatiojt, which dissolves out the sodium carbon- 
ate. From this solution it is recovered by evaporation. 
The undissolved portion is called the soda residua. 

During the black-ash process, some of the limestone is 
decomposed by the high temperature, forming calcium 
oxid and carbon dioxid, and when the black -ash is lixivia- 
ted the calcium oxid acts upon some of the sodium 
carbonate, forming sodium hydroxid, thus: 

CaO + Na 2 CO s -f H 2 = CaCO s + 2 NaOH. 

For this reason the crude soda, formed by the Leblanc 
process, always contains more or less sodium hydroxid. 

During the different processes a number of by-products 
are obtained. The hydrochloric acid, obtained in the salt- 
cake process, is employed in various ways, and the sulfur 
is recovered from the calcium sulfid in the soda residue. 

252. The Ammonia-Soda Process. This process is based 
upon the fact, that when sodium chlorid is acted upon by 
acid ammonium carbonate in solution, there is formed 
ammonium chlorid, and acid sodium carbonate, thus: 

NaCl + NH 4 HC0 3 = NH 4 C1 -f NaHCO„. 



206 LECTURE NOTES 

The acid sodium carbonate is then changed into the 
normal carbonate by ignition, thus: 

2NaHC0 3 = Na 2 C0 3 + H 2 + C0 2 . 
In actual practice the salt solution is saturated with 
ammonia, and this solution then charged with carbon 
dioxid under pressure, thus: 

NaCl -f NH 3 + H a O + C0 2 = NH 4 C1 + NaHC0 3 . 

This process is very simple from the chemical stand- 
point, but presented so many technical difficulties, that 
although the process was patented as early as 1838, it has 
only become prominent within a comparatively few years. 
The difficulties have been completely overcome, and largely 
through the efforts of a Belgian manufacturer named Sol- 
vay, and so the process is often called the Solvay process. 

At the present time considerably more soda is manu- 
factured by this, than by the Leblanc process. It yields 
a purer product at a less cost, and, with care, very little 
waste occurs. The raw material is cheap, and there are 
few by-products. 

253. Other Soda Processes. If the mineral cryolite, 
A1F 3 , 3 NaF, is fused with limestone, there is formed 
calcium fluorid and sodium aluminate, Na 3 A10 3 . The latter 
is soluble in water, and if the solution is charged with 
carbon dioxid, sodium carbonate is formed, thus: 

2 Na 3 M0 3 -[- 3 H 2 + 3 CO, = 2 Al(OH) 8 + 3 Na,CO g - 

An electrolytic process has lately been devised, which 
threatens to supersede all others. 

By this method sodium chlorid is first decomposed, as 
given under sodium hydroxid (243). The hydroxid is 
then charged with carbon dioxid, forming sodium car- 
bonate. 



LECTURE NOTES 207 

254. Acid Sodium Carbonate, NaHC0 3 . This compound 
is known commercially as bicarbonate of soda, and as 
baking soda. It is produced in enormous quantities in 
the ammonia-soda process, and may also be produced by 
the action of carbon dioxid upon the normal carbonate, 
thus: 

Na 2 C0 3 -f C0 2 + H 2 = 2 NaHCO,. 

Acid sodium carbonate is soluble in water, but much 
less so than the normal carbonate. If it is heated, or if a 
solution is boiled for a long time, it decomposes, forming 
the normal carbonate and carbon dioxid. 

255. Sodium Tetraborate, or Borax, Na 2 B 4 7 . This 
important salt occurs as the mineral tincal, or borax, in 
many localities, especially in Thibet, California, and 
Nevada. 

It is obtained from this source by recrystallization. It 
is also formed by boiling or fusing boric acid with sodium 
carbonate, thus: 

4 H a BO s + Na^CO, = Na,B,0, + 6 H,0 + CO,. 

Borax is easily soluble in water, from which it crystal- 
lizes in large, transparent prisms, with ten molecules of 
water of crystallization, Na 2 B 4 7 , 10 H 2 0. When the 
crystals are heated they swell up, lose their water of 
crystallization, and, on further heating, melt to a clear 
glass. This fused borax dissolves many of the metallic 
oxids, which impart characteristic colors to the glass. It 
is, therefore, used in blowpipe analysis for making the 
so-called borax beads for the detection of certain metals. 

Borax is used in the arts for soldering and welding, for 
glazing earthenware, and for many other purposes. It 
is also an antiseptic and so used in medicine. 



208 LECTURE NOTES 

Potassium. Symbol, K. (Kalium). Atomic Weight. 39.1. 

256. History and Occurrence. Some of the compounds 
of potassium have been known since the very earliest 
times, but the metal itself was first obtained by Davy, in 

1807. 

Potassium is the seventh most widely distributed ele- 
ment known. It constitutes about 2.4 per cent, of the 
earth's crust. It does not occur free in nature, but its 
compounds are very numerous and important. 

Among the most important natural compounds of potas- 
sium are sylvite (KC1), and carnallite (KC1, MgCl 2 , 6 H a O), 
both of which are found in the salt deposits in various 
parts of the earth, especially at Stassfurt in Germany. 
These minerals are the principal sources of potassium and 
its compounds. 

Potassium is also found in small quantities in sea- water, 
and in salt lakes and springs. It is found in all fruitful 
soils, and from these it is taken up by all plants, where 
it occurs as the salt of some organic acid. When the 
plants are burned, the potassium is found in the ash as 
the carbonate. From the plants it gets into the animal 
body, where it forms a small but important constituent. 

257. Preparation. Potassium was first prepared by 
Davy by electrolysis of potassium hydroxid. Until re- 
cently, all the potassium was obtained by strongly ignit- 
ing a mixture of potassium carbonate and carbon in an 
iron retort. This mixture is best obtained by igniting 
crude tartar, or argol, which is composed principally of 
acid potassium tartrate, HK(C 4 H 4 6 ), thus: 

2 HK(C 4 H 4 6 ) = K 8 C0 8 + 3 C -f 5 H 2 + 4 CO. 



LECTURE NOTES 209 

The final action is similar to that given under sodium, 
thus: 

K 2 CO s +2C = 2K-r-3CO. 

At the high temperature at which this action takes 
place, the potassium combines with the carbon monoxid, 
forming a black compound, K 6 (CO) 6 , which is very explo- 
sive. In order to avoid danger, the vapors from the 
retort are made to pass into a condenser of peculiar con- 
struction, where they are cooled rapidly and so the for- 
mation of the explosive compound is prevented. 

Potassium is also obtained by the electrolysis of potas- 
sium hydroxid by Castner's process, as already given 
under sodium (242). 

258. Properties and Uses. Pure potassium is a silver- 
white metal, with a bright metallic luster. It oxidizes 
very quickly in moist air, and decomposes water so ener- 
getically that the heat of the action is sufficient to set fire 
to the liberated hydrogen. It is, therefore, more electro- 
positive than sodium. 

At 0° it is hard and brittle, but at 15° it becomes soft, 
like wax. It melts at 62.5° and boils at 667°, sfivinsr a 



g ] 



splendid green vapor. If cooled in an atmosphere of hy- 
drogen it forms regular crystals. 

It is a somewhat more powerful reducing agent than 
sodium, but is very little used for this purpose, sodium 
being very much cheaper. 

Potassium and all of its compounds, give a violet color 
to the non-luminous gas flame. 

COMPOUNDS OF POTASSIUM. 

259. Potassium Hydroxid, KOH. This important com- 



210 LECTURE NOTES 

pound, which is also known as caustic potash, is exactly 
analogous to sodium hydroxid, and may be obtained in a 
similar way. 

Potassium hydroxid is a white crystalline solid. It is 
extremely deliquescent, and dissolves very easily in water 
with evolution of heat. It readily absorbs carbon dioxid, 
and is much used for this purpose in the laboratory. It 
is a powerful caustic, destroying both animal and vegeta- 
ble substances. 

Its most important commercial use is in the manufac- 
ture of soft soap. For most purposes, the more common 
and cheaper sodium hydroxid is used. 

260. The Halogen Compounds of Potassium. Potassium 
forms a compound with each of the halogen elements. 
They are all soluble in water and crystallize in cubes. 
The compound with fluorin has no commercial value. 

Potassium chlorid, KC1, occurs in sea- water and in 
springs. It also occurs as the mineral sylvite, and, com- 
bined with magnesium chlorid, as carnallite, in the Stass- 
furt and other salt deposits. It is mostly obtained from 
carnallite by solution and recrystallization. 

Potassium chlorid is used for the preparation of many 
of the other salts of potassium, and also in the preparation 
of fertilizers. 

Potassium bromid, KBr, is a valuable medicine and 
much used in nervous diseases. 

Potassium iodid, KI, is also used in medicine. 

261. Potassium Chlorate, KC10 3 . This important com- 
pound was probably known to the early alchemists. It 
is formed by leading chlorin into a solution of potassium 
hydroxid, thus: 



LECTURE NOTES 211 

6 KOH + 3 Cl 2 = K01O 3 + 5 KC1 + 3 H 2 0. 

The potassium chlorate being less soluble than the 
chlorid, crystallizes out, leaving the latter in the solution. 
In the manufacture of this substance, in order to avoid 
the loss of so much potassium in the form of chlorid, 
calcium hydroxid is used instead of potassium hydroxid. 
This forms calcium chlorate, which is then treated with 
potassium chlorid, forming potassium chlorate and calcium 
chlorid, thus: 

Ca(C10 3 ) 2 + 2 KC1 = 2 KC10 3 + CaCl 2 . 

The potassium chlorate is then separated by crystalli- 
zation, as in the process given above. 

A new method, based upon the fact that potassium 
chlorate is produced by the electrolysis of an alkaline 
solution of potassium chlorid, seems likely to replace the 
older methods in the near future. [Compare 244.] 

Potassium chlorate crystallizes in white tabular crystals, 
which often show iridescent colors. It melts at 359° and 
at a slightly higher temperature begins to decompose into 
potassium perchlorate, KCIO^, potassium chlorid, and 
oxygen, thus: 

2 KC10 3 = KC10 4 -f KC1 + 2 . 

At a still higher temperature the perchlorate decom- 
poses into the chlorid and oxygen, thus: 

KC10 4 = KC1 + 2 2 . 

When mixed with red phosphorus, or sulfur, or certain 
sulfids, it explodes violently on being heated or struck. 

Potassium chlorate is a powerful oxidizing agent and is 
largely used in pyrotechny and in the manufacture of 
matches (144). 



212 LECTURE NOTES 

262. Potassium Sulfate, K 2 S0 4 . This compound is 
found in the minerals kainite, (K 2 S0 4 , MgS0 4 , MgCl 2 , 6 H 2 0, ) 
and other similar minerals, which occur in the salt depos- 
its at Stassfurt. It is either obtained from these, or made 
by the action of sulfuric acid on potassium chlorid. 

It forms small colorless crystals which dissolve in water 
with difficulty, and contain no water of crystallization. 
It is used in the manufacture of alum, and for fertilizers. 

263. Potassium Nitrate, KN0 3 . This compound, which 
is known commercially as saltpeter, or niter, has been 
known since very early times. It is the most valuable, 
commercially, of all the potassium salts. 

Potassium nitrate occurs as an efflorescence upon the 
surface of the earth in various hot countries, particularly 
in India, Ceylon, Egypt, etc. It is formed as the result 
of the oxidation of organic matter containing nitrogen, 
in the presence of the potassium salts in the soil. This 
is brought about by the aid of certain specific organisms, 
or microbes, without which the action never takes place. 
This natural process is sometimes carried on artificially 
on the so-called saltpeter plantations. 

Potassium nitrate is also produced in large quantities 
by boiling a solution containing equivalent amounts of 
potassium chlorid and sodium nitrate, thus: 

NaN0 3 + KC1 = NaCl -f KNO,. 

Potassium nitrate forms large prismatic crystals with- 
out water of crystallization. It melts at 339°, and at a 
high temperature loses an atom of oxygen and forms 
potassium nitrite, KN0 2 . At a high temperature it is, 
therefore, an oxidizing agent. It is principally used in 
pyrotechny, and in the manufacture of gunpowder. 



LECTURE NOTES 213 

264. Gunpowder. This very important commercial 
product is a mixture of potassium nitrate, charcoal, and 
sulfur, in proportions which vary somewhat according to 
the use for which it is intended. The Chinese were prob- 
ably the first to make gunpowder, which they have used 
in making fireworks since very early times. Its use in 
gunnery dates from the early part of the 14th century. 

The most effective gunpowder is found to be that which 
contains about two molecules of saltpeter, to three atoms 
of carbon, and one of sulfur. It was at first believed that 
the chemical action, which takes place when this mixture 
is burned, could be represented by the following simple 
equation: 

2 KN0 3 -|- 3 C + S = K 2 S + N 2 + C0 2 . 

Careful analysis of both the solid and gaseous products 
of the combustion, shows that the decomposition is in fact 
very complex. The solid part, which consists of the resi- 
due left in the gun, together with the smoke, contains 
very little potassium sulfid, the principal constituents 
being potassium sulfate, carbonate, and thiosulfate, with 
some unchanged saltpeter, and carbon, while the gases 
contain carbon monoxid, hydrogen sulfid, and hydrogen, 
in addition to the nitrogen and carbon dioxid. The hy- 
drogen comes from the charcoal, which always contains 
more or less hydrogen and oxygen, in the form of hydro- 
carbons. 

The United States military gunpowder contains 76 per 
cent, saltpeter, 14 per cent, carbon, and 10 per cent, sul- 
fur. Sporting powder contains two per cent, more salt- 
peter in place of carbon. Blasting powder contains 62 
per cent, saltpeter, 18 per cent, carbon, and 20 per cent, 
sulfur. 



214 LECTURE NOTES 

Gunpowder takes fire at about 300°, and the tempera- 
ture of the combustion is about 2200°. The amount of 
gas formed by the explosion is about 43 per cent, of the 
gunpowder, by weight. The expansion of this gas caused 
by the heat of explosion, under the most favorable circum- 
stances, gives an initial pressure of about 6400 atmos- 
pheres, or 96000 pounds per square inch. 

265. Potassium Carbonate, K 2 C0 3 . This salt was form- 
erly obtained by leaching the ashes of wood and other land 
plants, and was known by the name of potashes. The 
crude product is dark colored, but becomes white by crys- 
tallization, and ignition to burn off organic matter, and 
is then called pearl-ash. 

It is now manufactured by a process exactly similar to 
the Leblanc soda process. 

Potassium carbonate forms long prismatic crystals with 
three molecules of water, K 2 C0 3 , 3 H 2 0. It is deliques- 
cent and very soluble in water. It melts at 879°, and 
on further heating loses a little carbon dioxid. 

It is used for making soft soap and in the manufac- 
ture of glass. 

266. Potassium Cyanid, KCN. This compound is com- 
monly prepared by heating potassium ferrocyanid, thus: 

K 4 Fe(CN) 6 = 4 KCN -f FeC 2 + N 2 . 

Potassium cyanid is a white solid. It is very soluble 
in water and is decomposed by the weakest acids, such 
as carbonic acid. It melts easily and in this condition 
takes on oxygen forming potassium cyanate, KOCN, and 
so is a reducing agent, and used as such in the laboratory. 
It is exceedingly poisonous. 

Its principal use, commercially, is for recovering gold 



LECTURE NOTES 215 

from low grade ores. A mixture of potassium and sodium 
cyanids is used for this purpose, which is made by heat- 
ing potassium ferrocyanid with sodium, thus: 

K 4 Fe(CN) 6 + 2 Na = 4 KCN + 2 NaCN + Fe. 

267. Rubidium and Caesium. These two very rare 
elements have a very peculiar interest attached to them, 
inasmuch as both were discovered by Bunsen in the waters 
of Durkheim, by means of the spectroscope. 

They are found in minute quantities in a number of 
mineral waters and in a few minerals. They very much 
resemble potassium and, by ordinary analytical reactions, 
they cannot be distinguished from it, or from each other. 
They are more electro-positive than potassium, caesium 
being the most electro-positive element known. 

Rubidium imparts a redder color to the non -luminous 
gas flame than does potassium, and the caesium flame is 
still redder than that of rubidium. 

268. The Compound Radical Ammonium, NH 4 . When 
ammonia gas, NH 3 , is brought into contact with any 
acid, it forms by addition what is called an ammonium 
compound (127). All such compounds contain the com- 
pound radical, NH 4 , which is called ammonium, and which 
may be regarded as a metallic group, since, in its com- 
pounds, it resembles the alkali-metals. The ammonium 
compounds, therefore, may also be regarded as having 
been derived from acids, by replacing hydrogen in the 
acid by means of the compound radical ammonium. Am- 
monium itself has never been isolated, but its existence 
in compounds is unquestionable. 

The ammonium compounds are all easily volatile, and 
since they are in other respects analogous to those of the 
alkali-metals, ammonium is often called the volatile alkali. 



216 LECTURE NOTES 

If sodium amalgam is placed in a concentrated solution 
of ammonium chlorid, a very curious action takes place. 
The mercury increases enormously in bulk, and forms a 
light, soft mass, which is called ammonium amalgam, 
and which soon decomposes into mercury, hydrogen and 
ammonia. 

It is believed by some that the so-called ammonium 
amalgam is not a true amalgam, but a sort of solution of 
ammonia and hydrogen in mercury. In opposition to 
this belief the fact remains that neither ammonia nor 
hydrogen is in the least degree soluble in mercury, and 
that in other ways the substance does resemble a true 
amalgam. 

THE AMMONIUM COMPOUNDS. 

269. Ammonium Hydroxid, NH 4 OH. The existence 
of a compound having this composition is not yet fully 
proven. If ammonia gas is dissolved in water, there is 
formed a strongly alkaline solution, which neutralizes 
acids, forms hydroxids of the heavy metals from solutions 
of their salts, and, in general, deports itself like a base 
of the alkali-metals. If a solution of ammonia is heated, 
ammonia passes off, and if the solution is boiled for some 
time, all the ammonia will pass off, which shows that if 
ammonium hydroxid exists it is very unstable. 

That there is a weak chemical union of ammonia and 
water when the former dissolves in the latter, seems highly 
probable from the fact, that while ammonium hydroxid 
cannot be isolated, there are several organic compounds 
which are derivatives of it, and which resemble it very 
closely. The belief in the existence of this compound in 
solution, is, therefore, quite general, and certainly helps 



LECTURE NOTES 217 

us to a clearer understanding of the other ammonium 
compounds. 

270. Ammonium Chlorid, NH 4 C1. This important com- 
pound occurs in small quantities in nature in volcanic 
districts, and is known commercially as sal-ammoniac. 

It is prepared by leading ammonia gas, obtained from 
the ammoniacal liquor of the gas works (125), into hydro- 
chloric acid. The solution of ammonium chlorid thus 
obtained, is evaporated to dryness, and the salt purified 
by sublimation. 

Ammonium chlorid forms white, fern-like crystals, 
which consist of groups of minute octahedrons. It is 
very soluble in water, the solution being accompanied by 
a considerable loss of heat. On boiling a solution of am- 
monium chlorid, a partial dissociation of the salt takes 
place. Some ammonia passes off. and the solution becomes 
slightly acid. 

If ammonium chlorid is heated under ordinary atmos- 
pheric pressure, it sublimes without melting, and, at the 
same time, dissociates into ammonia and hydrochloric 
acid. The dissociation is complete at 350°, but on cool- 
ing the two gases recombine forming ammonium chlorid. 

Ammonium chlorid is an important reagent in the lab- 
oratory where it is also used as a convenient source of 
ammonia. It is also used in the arts, and in medicine 

271. The Ammonium Sulfids. If an aqueous solution 
of ammonia (ammotiium hydroxid) is saturated with hy- 
drogen sulfid, there is formed principally ammonium 
hydrosulfid, NH 4 SH. 

If this solution is mixed with an equal volume of ammo- 



218 LECTURE NOTES 

nium hydroxid there is formed colorless ammonium sulfid, 
(NH 4 ) 2 S, thus: 

NH 4 SH -f NH 4 OH = (NH 4 ) 2 S + H 2 0. 

If sulfur is dissolved in either of the above solutions, 
there is formed a yellow ammonium sulfid, or ammonium 
poly sulfid, (NH 4 ) 2 S X . There are several of these polysul- 
fids, the number of sulfur atoms varying from two to 
seven. They all contain more or less of some other unim- 
portant compounds which are closely related to the above. 

All of the ammonium sulfids are easily decomposed by 
heat, and, in solution, are more or less completely disso- 
ciated by standing, especially in contact with the air and 
in the light. 

They are all used in analytical chemistry for precipita- 
ting the sulfids of the metals, and for separating them. 

272. Ammonium Sulfate, (NH 4 ) 2 S0 4 . This compound 
is found in certain volcanic districts, and is made bypass- 
ing ammonia gas into sulfuric acid, after which it is 
purified by recrystallization. 

It forms large colorless crystals without water of crys- 
tallization. It melts at 140° and decomposes at a higher 
temperature. 

It is used in the manufacture of other ammonium salts 
and as a fertilizer. 

273. Ammonium Nitrate, NH 4 N0 3 . This compound is 
formed by leading ammonia into nitric acid. It crystal- 
lizes from a solution in different forms according to the 
temperature at which the crystallization takes place. 

If ammonium nitrate is heated, it melts at 159°, and at 
170° it decomposes into nitrous oxid, N 2 0, and water(129), 



LECTURE NOTES 219 

a small portion of it subliming unchanged. If the sub- 
stance is heated to about 240°, it decomposes with violent 
explosion into nitrogen, nitric oxid (NO), and water. 

274. Ammonium Sodium Hydrogen Phosphate, 

NH 4 NaHP0 4 . This important compound, which is also 
called microcosmic salt, was known to the alchemists, who 
first obtained it from urine. It is also found in guano. 

It can be formed by dissolving sodium and ammonium 
phosphates in hot water, and allowing the solution to 
cool. It is also obtained by the action of a strong solu- 
tion of sodium phosphate on ammonium chlorid, thus: 
Na HPO, + NH CI = NH NaHPO, -I- NaCl. 

2 4 ■ 4 4 4 ' 

It forms large transparent prismatic crystals having 
four molecules of water of crystallization. When heated 
it melts easily, loses water and ammonia, and, if strongly 
heated, forms a clear glass of sodium metaphosphate, 
NaP0 3 . This dissolves many of the metallic oxids, which 
impart characteristic colors to the glass. It is, therefore, 
much used in blowpipe analysis. 

The ordinary phosphates of ammonium resemble some- 
what the corresponding compounds of sodium, but are 
not very important. 

275. Ammonium Carbonate, (NH 4 ) 2 C0 3 . When ammo- 
nium sulfate and ground chalk (CaC0 3 ) are heated in an 
iron retort, commercial ammonium carbonate is formed. 
This substance, which is also called sal volatile, is a mix- 
ture of acid ammonium carbonate, NH 4 HC0 3 , and ammo- 
nium carbamate, NH 4 C0 2 NH 2 . When a solution of this 
substance is treated with ammonia gas the normal car- 
bonate is formed. The ammonia acts upon the carbonate, 
thus: 

NH 4 HCO, + NH 3 = (NH 4 ) 2 C0 3 . 



220 LECTURE NOTES 

The water decomposes the carbamate, thus : 

NH 4 C0 2 NH 2 + H 2 = (NH 4 ),CO,. 

Ammonium carbonate is a white crystalline substance, 
which gives off ammonia when exposed to the air, and 
forms the acid carbonate. If the normal carbonate is 
heated to 60°, it is entirely decomposed into ammonia, 
carbon dioxid, and water. 

276. The Copper Group. If we examine the table on 
page 192, we shall see that in addition to the five alkali- 
metals, Group I contains the metals copper, silver, and 
gold. These three metals differ considerably from the 
alkali-metals, but they form a series of compounds in 
which the metal is monovalent. They also somewhat 
resemble the connecting metals of G-roup VIII in the 
middle of the long series. We shall also see that in the 
arrangement of the table the second metal in the long 
periods is placed at the right in each group. They form 
what is called division B of each group, the metals at the 
left forming division A. 

These two divisions quite closely resemble each other, 
but the metals of each division resemble one another, 
much more closely than they do those of the other 
division. 

Copper. Symbol, Cu. (Cuprum). Atomic Weight, 63.6. 

277. History and Occurrence. Copper has been known 
since the very earliest times, and was probably the first 
metal to be used by man. The names by which it was 
first known seem to have been used indiscriminately for 
copper, and the alloys brass and bronze. Although native 
copper is found in Palestine and the surrounding coun- 
tries, its principal source was the island of Cyprus. It 



LECTURE NOTES 221 

was therefore called by the Romans aes cyprium, or 
cyprium, and later cuprum. For this reason copper was 
regarded as sacred to Venus, and in the writings of the 
alchemists it is called Venus and given the symbol, 9? 
supposed to represent a mirror, which was one of the 
attributes of the goddess. 

The alchemists knew the fact, that when iron was placed 
in a solution of a copper salt metallic copper was ob- 
tained, and this action was regarded by them as a trans- 
mutation of iron into copper. 

Copper occurs in the native state in all parts of the 
world. Perhaps its most important locality is the Lake 
Superior region, where it occurs in enormous masses. It 
is also found in a large number of minerals, the more 
important of which are cuprite, (Cu 2 0), chalcocite (Cu a S), 
bornite (3 Cu 2 S,Fe 2 S s ), chalcopyrite (Cu 2 S,Fe 2 S 3 ), malachite 
(CuC0 3 ,Cu(OH) 2 ), and azurite (2 CuC0 3 ,Cu(OH) 2 ). 

278. Preparation. The Dry Process. The methods by 
which copper is obtained from its ores vary considerably, 
and depend upon the nature of the ore. 

If the ore is the oxid, or some compound easily re- 
duced to the oxid, like the carbonate, the metal is obtained 
by reducing the ore in a blast furnace by means of coal 
or coke, thus: 

Cu 3 4- C = 2 Cu -f CO. 

By far the greater portion of copper ore contains more 
or less sulfur, in which case the metal is obtained by the 
English method. This consists of a number of distinct 
processes. 

First, the ores, which consist of sulfids of copper and 
iron, together with some silica, and perhaps other impu- 



222 LECTURE NOTES 

rities, are strongly heated in a reverberatory furnace. By 
this action they are partially oxidized, the sulfur forming 
sulfur dioxid. This is called calcining the ores. 

The second step is to fuse the calcined ores, by which 
means the oxids of copper are made to react upon the 
iron (ferrous) sulfid, thus: 

Cu 2 + FeS = Cu 2 S + FeO, and 
6 CuO + 4 FeS = 3 Cu 2 S -f 4 FeO + S0 2 . 

The iron oxid combines with the silica in the ore to 
form a fusible silicate of iron, called slag, and is drawn 
off. If the ore does not contain sufficient silica, some of 
the metal-slag from a later process is added. The cu- 
prous sulfid, with some unchanged iron sulfid, remains 
behind as a regulus, and is called coarse metal. 

The coarse metal, which contains about 35 per cent, of 
copper in the form of sulfid, is then calcined, which re- 
moves some more of the sulfur as sulfur dioxid, and 
partially oxidizes the metal. 

The calcined mass is then fused with enough refinery- 
slag (obtained in a later operation), to combine with the 
remaining' iron, which forms the so-called metal-slag. 
The regulus then consists of nearly pure cuprous sulfid, 
and is called white-metal or fine-metal ', and contains about 
75 per cent of copper. 

The white metal is then heated again in a reverberatory 
furnace. Here again some of the sulfid is changed to oxid 
after which the two react upon each other, thus: 

2 Cu,0 -|- Cu 2 S = 6Cu + S0 2 . 

The copper thus obtained is not yet quite pure and is 
called blister -copper, because of its appearance. 



LECTURE NOTES 223 

The final process is the refining of the blister-copper. 
For this purpose it is melted in a reverberatory furnace, 
where the impurities, lead, arsenic, etc., are oxidized, and 
either volatilize, or are removed as slag. The molten 
copper itself, having become somewhat oxidized, is then 
covered with a thin layer of coal, and stirred with poles of 
green wood, until it attains the proper degree of toughness. 

279. Preparation. The Wet Process. In the manu- 
facture of sulfuric acid, enormous quantities of pyrite are 
burned for the preparation of sulfur dioxid. This pyrite 
contains, on the average, about 3 per cent, of copper, 
which amount is too small to recover by the methods just 
given. 

The burned ore, known commercially as blue-billy, is 
therefore ground, mixed with from 12 to 15 per cent, of 
rock-salt, and calcined. This converts the copper into 
copper chlorid. The mass is then lixiviated with water, 
which dissolves the copper chlorid. The copper is then 
precipitated by means of scrap iron, and refined. 

280. The Refining of Copper by Electrolysis. The crude 
copper, obtained by the above methods, often contains 
enough gold and silver to make it profitable to extract 
them, and this is now quite largely done by electrolysis. 

By this method the crude copper (blister copper) is first 
cast into plates, which are hung in long wooden tanks 
lined with lead, and form the positive pole, or anode, very 
thin sheets of pure copper being used for the negative 
pole, or cathode. The tanks are filled with a solution of 
copper sulfate, acidified with sulfuric acid. 

When a current of electricity is passed through the 
solution, the anode is dissolved, the copper being depos- 



224 LECTURE NOTES 

ited on the cathode, while the gold and silver, not being 
dissolved, fall to the bottom. Any other metals found 
in the crude copper, either remain in solution, or are 
precipitated with the gold and silver. With care the 
copper thus obtained is nearly chemically pure. 

281. Electro typing. The principle involved in the re- 
fining of copper by electrolysis, is the one employed in 
the very important process called electrotyping. 

Whatever may be the form of the cathode, whether a 
finely engraved plate, a page of type, or a figure of any 
kind, the deposit of copper upon it is found to be a perfect 
impression. 

. In the actual process, a reverse of the object is first 
made, either in plaster of Paris, or gutta-percha. This 
is then coated with graphite powder, so as to give con- 
ductivity to the surface, and is made the cathode in a 
copper sulfate solution, the anode being a plate of copper. 
An electric current of suitable strength is then passed 
through the solution, until the deposited film is of suffi- 
cient thickness. It is then taken out and the back filled 
with type-metal, or solder, so as to make it sufficiently 
rigid. 

282. Properties and Uses. Pure copper is a lustrous 
brownish-red metal. The native copper is often found 
crystallized in regular octahedrons. It is one of the 
toughest of the metals, and is very malleable and ductile, 
but when heated nearly to the melting point it becomes 
quite brittle. Next to silver it is the best conductor of 
heat and electricity. It has a specific gravity of 8.94, 
and melts at 1080°. 

Copper remains unchanged in perfectly dry air at the 



LECTURE NOTES 225 

ordinary temperature, but when heated it forms black 
copper oxid, CuO. In moist air it gradually becomes 
coated with a greenish basic carbonate. 

Copper is not ordinarily acted upon by dilute hydro- 
chloric or sulfuric acids, but in contact with a piece of 
platinum, or if moistened and exposed to the air, it slowly 
dissolves in these acids. It dissolves in hot concentrated 
sulfuric acid forming copper sulfate and sulfur dioxid 
(101). It dissolves easily in nitric acid forming copper 
nitrate and nitric oxid (131). 

Copper is very extensively used in making certain 
alloys, such as brass and bronze, which are of great com- 
mercial value. It is almost universally used for electric 
wiring, and all other electric connections. 

COMPOUNDS OF COPPER. 

283. Cuprous and Cupric Compounds. Copper forms 
two series of compounds, known as cuprous and cupric. 

In their constitution the cuprous compounds are exactly 
analogous to those of the alkali-metals, the copper being 
univalent. Cuprous oxid occurs in nature as cuprite, 
and is a fairly stable compound. All of the other cuprous 
compounds are very unstable, and when exposed to the 
air, are oxidized and form the corresponding cupric com- 
pounds. None of them are very important. Cuprous 
salts of the oxy-acids do not exist. 

The cupric compounds, which include all the ordinary 
compounds of copper, are all quite stable, and the copper 
in them is bivalent. They resemble in many respects the 
corresponding compounds of the magnesium metals, and 
the compounds of iron and manganese in which the metals 
are bivalent. 



226 LECTURE NOTES 

284. Copper Oxid, CuO. This compound occurs in na- 
ture as the rather rare mineral melaconite. It can be 
made by strongly heating copper in the air, or by gently 
igniting the nitrate, carbonate, or hydroxid. 

Copper oxid is a black, amorphous powder, which is 
hygroscopic. When strongly heated, it first cakes to- 
gether, and at a high temperature fuses, giving up a 
part of its oxygen. When heated in an atmosphere of 
hydrogen, or carbon monoxid, it is easily reduced to the 
metallic state, Reduction also takes place when it is 
heated with any compound of carbon, and for this reason 
it is largely used in the ultimate analysis of organic com- 
pounds. 

Copper oxid, and its corresponding hydroxid, Cu(OH) 2 , 
dissolves in ammonium hydroxid forming a deep blue liquid. 
This solution dissolves cellulose (cotton, wool, filter paper, 
etc.), and is known as Schweitzer's reagent. 

285. Copper Sulfate, CuS0 4 . This important salt can 
be formed by dissolving copper in concentrated sulfuric 
acid. It is made on a commercial scale, by carefully 
roasting ores containing copper sulfid, the latter being 
thus oxidized to the sulfate. The residue is lixiviated with 
water and the copper sulfate obtained from the solution 
by crystallization. 

Copper sulfate forms large blue crystals which contain 
five molecules of water, CuS0 4 , 5 H 2 0, and are known 
commercially as blue vitriol. These crystals lose four 
molecules of water at 100°, and the fifth at about 240°. 
The anhydrous salt is white and very hygroscopic, and is 
used to detect and remove small quantities of water from 
organic liquids. 



LECTURE NOTES 227 

Copper sulfate forms with the alkaline sulfates, a 
series of double salts, which have the general formula 
CuS0 4 , M' 2 S0 4 , 6 H a O. These are isomorphous with the 
so-called monoclinic double salts formed by the metals of 
the magnesium group. (Compare 328). 

Copper sulfate is used extensively in calico printing, in 
electrotyping, and in the preparation of the pigments 
known as Scheele's green and Paris green. 

286. Other Salts of Copper. The other salts of copper 
are not of particular importance. They are formed by 
the general methods already given (234). 

Copper Chlorid, CuCl 2 ,2 H 2 0, crystallizes in fine green 
prisms. On strongly heating, these lose water and part 
of their chlorin, and form cuprous chlorid, Cu 2 Cl 2 . 

Copper Nitrate, Cu(N0 3 ) 2 , 3 H 2 0, forms dark green 
deliquescent crystals. This salt is easily decomposed, and 
so is a strong oxidizing agent. It is used for this purpose 
in calico printing and dyeing. 

Normal Copper Carbonate, CuC0 3 , is not known. There 
are two basic carbonates, which occur in nature as mala- 
chite and azurite. Malachite can be formed artificially by 
precipitating a solution of a copper salt with sodium car- 
bonate. It is a green compound, having the composition 
CuC0 3 , Cu(OH) 2 . It has the same composition as the so- 
called verdigris, or copper rust. Azurite has the composi- 
tion 2 CuC0 3 , Cu(OH) 2 , and forms beautiful dark azure- 
blue crystals. 

Copper salts form a number of interesting compounds 
with ammonia. Very little is at present known as to 
their exact constitution, though they are generally re- 
garded as compound molecules formed by addition. They 
all dissolve in water forming a deep blue solution. 



228 LECTURE NOTES 

Silver. Symbol, Ag. (Argentum). Atomic Weight, 107.9. 

287. History and Occurrence. Silver has been known 
since the earliest times, and the various names given to 
the metal by the ancients, referred to its bright white 
color. The alchemists called it Luna, and usually repre- 
sented it by the symbol for the moon, which was the cres- 
cent, 3. 

Silver occurs in the native state in many parts of the 
earth, and occasionally in quite large masses. Its most 
important ores are argentite, Ag 2 S, pyrargyrite or ruby 
silver, Ag 3 AsS 3 , proustite, AgSbS 3 , and chlorargyrite or 
horn silver, AgCl. Of these the argentite is by far the 
most important. The common lead ore galenite, PbS, 
nearly always contains some silver, the amount varying 
from a mere trace to one per cent, or even more. When 
the argentiferous galena, as is called, is mined for silver, 
it ranks as a silver ore, and this constitutes one of the 
main supplies of silver. 

288. Preparation. The Dry Process. The methods of 
extracting silver from its ores are quite complicated, and 
the one employed depends upon the amount of silver in 
the ore, and the conditions of locality. It is only from 
the richest ores that the silver can be obtained directly. 
The most important of all the methods is the one which 
depends upon the formation of an alloy of silver and lead, 
and the subsequent separation of the two metals. 

When silver ores are melted with metallic lead, or with 
an ore which yields metallic lead, an alloy of silver and 
lead is formed. When the alloy is rich in silver the 
metals are separated by the process called cupellation. 
This consists in heating the metals in a reverberatory 



LECTURE NOTES 229 

furnace, in a shallow dish made of bone ash, or clay, and 
called a cupel. A strong blast of air is thrown upon the 
surface of the molten metals, which oxidizes the lead, but 
does not change the silver. The melted lead oxid is 
drawn off and the pure silver remains. 

When the alloy is too poor in silver to be directly cu- 
pelled, it is. first submitted to an enriching process, of 
which there are two, known as the Pattinson process, and 
the Parkes's process. 

The Pattinson process depends upon the fact, that an 
alloy of silver and lead melts at a lower temperature than 
pure lead, so that when the molten alloy is cooled, the 
crystals which are first formed are nearly pure lead. The 
crystals are removed by means of a perforated iron ladle, 
and the operation continued until a definite proportion 
has been removed, which proportion depends upon the 
amount of silver originally present. The remaining lead 
is then removed from the rich silver alloy by cupellation. 

The Parties' s process depends upon the fact, that when 
zinc is added to a molten alloy of silver and lead, it alloys 
with the silver, and leaves the lead nearly pure. The 
zinc-silver alloy is lighter, and solidifies at a lower tem- 
perature than the lead. In the actual process the silver- 
lead alloy is melted, and an amount of zinc, equal to about 
twenty-five times the weight of the silver, added, and 
thoroughly stirred into the molten mass. When the mix- 
ture cools, the zinc-silver alloy solidifies and rises to the 
top, where it is skimmed off with a perforated iron ladle. 
The skimmings are then carefully heated, when most of 
the adhering lead runs away. The zinc is removed from 
the silver by distillation, after which the silver is purified 
by cupellation. 



230 LECTURE NOTES 

289. Preparation. The Wet Processes. There are 
several wet processes in operation at the present time. 
The process which is carried on in countries where fuel is 
scarce, is called the Amalgamation process, and is the one 
commonly used in Mexico and South America. This pro- 
cess depends upon the fact, that certain compounds of 
silver, such as the chlorid, are decomposed by mercury 
forming mercurous chlorid, the reduced silver forming an 
amalgam with the excess of the mercury. 

In the process, the ore, mixed with water, is ground to 
a fine powder in a mill operated by horse or mule power. 
The muddy mass thus obtained is spread upon a paved 
space, and thoroughly mixed with from 3 to 5 per cent, of 
sodium chlorid, the mixing being done by the treading of 
mules. Mercury is then added, together with a substance 
called magistral, which is obtained by roasting pyrite, 
and consists of a mixture of copper and iron sulfates and 
oxids. More mercury is added from time to time, as is 
found necessary, the whole operation lasting for several 
days. The mass is then carefully washed, the amalgam 
filtered through canvass bags and distilled. 

A modification of this process is employed in Nevada, 
and is called the Washoe process. 

Several other wet processes, made practical by local 
conditions, are in operation in different parts of the world. 

290. Properties. Pure silver is a lustrous white metal, 
which in the form of very thin foil transmits blue light. 
It has a specific gravity of 10.5 and melts at 960°. It is 
the best conductor for heat and electricity known, and, 
next to gold, the most malleable and ductile metal known. 
It remains unchanged in pure air, but is quickly tarnish- 
ed by air containing even traces of hydrogen sulfid. When 



LECTURE NOTES 231 

pure silver is melted, it absorbs oxygen, the maximum 
amount being about twenty-two times its own volume. 
When it cools the most of this oxygen is given off, and 
the bursting of the gas through the outer crust, ejects 
some of the still molten silver in curious forms. This 
phenomenon is known as the spitting of silver. A small 
amount of alloy prevents this absorption of oxygen. 

Silver combines directly with the halogens. It also dis- 
solves in hot concentrated sulfuric acid forming sulfur 
dioxid and silver sulfate. It dissolves easily in nitric acid, 
even when dilute, forming silver nitrate and nitric oxid. 

291. Uses. Silver, alloyed with copper, is very largely 
used for coinage, for ornamental purposes, and for silver 
plate, which is the name given to domestic articles made of 
silver. English silver coins contain 92. 5 per cent, of silver, 
while those of most other countries contain but 90 per 
cent. The term finen ess, as applied to silver, has reference 
to the number of parts of pure silver in 1000. English 
coin has, therefore, a fineness of 925. Silver plate has 
usually a fineness of 950, while silver jewelry often has a 
fineness of only 800. 

Silver plating is the process of covering articles made 
of some cheap metal or alloy, with silver. The plating 
solution is made by dissolving silver cyanid in potassium 
cyanid. The process is exactly analogous to that em- 
ployed in electrotyping (281). 

A number of organic compounds, such as aldehyde, 
tartaric acid, and milk sugar, will reduce silver salts in 
an alkaline solution. If a plate of clean glass is placed in 
the solution, silver will be deposited on the glass as a 
coherent mirror. Such mirrors reflect light very perfectly 
and are much used for optical purposes. 



232 LECTURE NOTES 



COMPOUNDS OF SILVER. 



292. The Halogen Compounds of Silver. Silver forms 
compounds with each of the halogens. The chlorid, 
bromid, and iodid, are much alike in both chemical and 
physical properties, and of considerable importance. 
When they are exposed to light, they first become purple, 
and finally become dark brown or black. For this reason 
they are much used in photography. The fluorid is very 
different from the other halogen salts, and is of little 
importance. 

Silver chlorid, AgCl, is found in nature as chlorargyrite, 
or horn silver. When a solution of any chlorid is mixed 
with a solution of a silver salt, silver chlorid is obtained 
as a white, curdy precipitate. It melts at about 450° 
forming a yellow liquid, which solidifies to a horn like 
mass. It is almost entirely insoluble in water and in 
concentrated acids, but dissolves very easily in ammonia. 

Silver bromid, AgBr, is analogous to the chlorid in the 
method of its preparation. It is a light yellow compound. 
It is exceeding sensitive to the action of light, and is the 
principal silver salt used in making photographic dry- 
plates. 

Silver iodid, Agl, is a yellow compound, which is more 
stable than either the chlorid or bromid. 

293. Silver Nitrate, AgN0 3 . This compound is made 
by dissolving silver in nitric acid. It forms large, white, 
tabular crystals, which melt at 200°, and on cooling forms 
a crystalline mass which is known commercially as lunar 
caustic. On being strongly heated, it first gives off some 
oxygen, forming silver nitrite, AgN0 2 , and at a still 
higher temperature decomposes into metallic silver and 
nitrogen peroxid. 



LECTURE NOTES 233 

If silver nitrate is brought in contact with organic 
matter, it blackens on exposure to the light, so that if 
the skin, or a piece of cloth, is moistened with a solution 
of this salt, it soon becomes black. It is on account of 
this property that it is much used for making indelible 
ink. 

It is easily soluble in water, and is the most common 
soluble salt of silver. It is much used in the laboratory, 
where it is the principal source of silver. 

294. Photography. The fact that silver salts blacken 
when exposed to the light, was known in the 17th centu- 
ry, and the application and development of this fact has 
resulted in the art of photography. 

In 1802, Wedgewood made prints of leaves, by placing 
them on paper moistened with silver nitrate and exposing 
them to the light. These pictures could not be kept in 
daylight, since no method of "fixing" them, or making 
them permanent, was known, and so they had no intrinsic 
value. Many experiments were tried to find some method 
of fixing the prints, but with little success until 1839, 
when the French painter Daguerre discovered the pro- 
cess which bears his name. 

In the daguerreotype process, a silver plate is exposed 
to the action of iodin vapor, thus forming a surface of sil- 
ver iodid. This is placed in a camera, and exposed for a 
short time to the action of the light. The exposed plate 
is then acted upon by the vapor of mercury, which brings 
out, or "developes," the picture, by depositing very fine 
globules of mercury on those parts of the plate which 
were acted upon by the light. The plate is then dipped 
in a solution of sodium thiosulfate which removes the un- 
changed silver salts, and thus "fixes" the plate, or renders 
it permanent. [Compare 111]. 



234 LECTURE NOTES 

This process underwent many changes and improve- 
ments, but after some years gave way to the process dis- 
covered by Archer in 1851, and known as the wet-collodion 
process. 

In this process a glass plate, covered with a film of 
iodized collodion, is dipped in a solution of silver nitrate, 
to render it sensitive to light, and then, while still wet, 
exposed in the camera. It is then acted upon by an acid 
solution of some reducing- agent, such as. ferrous sulfate, 
or pyrogallic acid, which reduces to metallic silver those 
silver salts which have been exposed to the light, and so 
developes the plate. The plate is then fixed by means of 
sodium thiosulfate, after which it is washed and dried. 

The necessity of having to prepare the wet-collodion 
plate immediately before using, made this process at 
times very inconvenient. To overcome this, dry plates 
were introduced, but the first of these, while convenient, 
were not very sensitive. These were gradually improved 
however, until in 1871, the bromid- gelatin process was 
discovered, a process which is substantially the one in use 
at the present time. 

In this process the plate is covered with an emulsion 
made of gelatin and silver bromid, with sometimes a little 
iodid, and dried. The emulsion is rendered more sensi- 
tive, or -'ripened," by heating some time before it is 
applied. These dry plates are exceedingly sensitive, the 
time required for exposure in the camera in a good light, 
being only a small fraction of a second. The plates are 
developed by an alkaline solution of pyrogallic acid, hy- 
droquinone, eikonogen,or some such reducing agent, and 
then fixed with sodium thiosulfate. 

From both the wet and the dry plates we obtain a 



LECTURE NOTES 235 

"negative, " that is, a plate on which the light portions 
of the object appear dark, while the dark parts are trans- 
parent. 

The finished photograph, or positive, is made by plac- 
ing the negative over paper coated with a film of albumen 
and silver chlorid, and exposing it to sunlight, a process 
called printing. The prints are then "toned" by being 
placed in a solution of gold chlorid, (other salts are some- 
times used), when the silver salts which have been acted 
upon by the light, reduce the gold chlorid, precipitating 
fine metallic gold. The pictures are then fixed by sodium 
thiosulfate, after which they are thoroughly washed and 
mounted on cardboard. 

Gold. Symbol, Au. (Aurum). Atomic Weight, 197.2. 

295. History and Occurrence. Gold was one of the first 
metals to be known and used by man. Its occurrence in 
the native state, together with its rich color and brilliant 
luster, has made it an object of value from the earliest 
times. 

The alchemists considered gold the most perfect of the 
metals and compared it to the sun. They also called it 
Sol, and gave it the symbol, Q, which was the symbol of 
the sun. The alchemists also believed gold to possess 
remarkable virtues, hence one of the great problems of 
alchemy was the transmutation of the base metals into 
gold. 

Gold is almost always found in the native state. It is 
never found perfectly pure but always alloyed with some 
silver, and occasionally with copper, or some of the plati- 
num metals. It is usually found in quartz veins inter- 
secting the oldest rocks, and to some extent in the rocks 



236 LECTURE NOTES 

themselves. By the disintegration of the rocks the gold 
gets into the gravels and soils, which form thus one of 
the chief sources of the metal. 

Although native gold occurs in all parts of the earth, 
it is only in a few localities that the quaDtity is sufficient 
to pay for extracting, some of the more important of these 
localities being Australia, South Africa, Russia, California 
and the Klondike in Alaska. 

G-old is also found combined, principally with tellurium, 
in the rare minerals sylvanite, (Au, Ag)Te 2 , and nagyagite, 
(Pb,Au) 2 (Te,S,Sb) 3 , and others. 

296. Preparation. Gold Mining. When gold is found 
in the soil or gravel, the process of extracting it is known 
as placer-mining. The gold is separated from the lighter 
mud and gravel by washing. The processes called panning, 
cradling, and hydraulic mining are all the same in 
principle. 

When the gold is found in quartz veins, the ore is first 
crushed in what is called a stamp mill, and then washed; 
or, the gold is extracted with mercury, with which it 
amalgamates, and the amalgam removed and distilled. 

Gold is extracted from auriferous pyrite by means of 
chlorin. The ore is first roasted to remove sulfur and 
arsenic, and after being moistened with water, is exposed 
to the action of chlorin. The gold chlorid thus formed, 
is extracted by lixiviating with water, and the gold pre- 
cipitated from the solution by means of ferrous sulfate, 
thus: 

AuCl 3 + 3 FeS0 4 = Au + FeCl 3 + Fe 2 (S0 4 ) 3 . 

A process, known as the cyanid process, and dependent 
upon the fact that finely divided gold is easily dissolved 



LECTURE NOTES 237 

in a solution of potassium cyanid in the presence of at- 
mospheric oxygen, is now employed on a very large scale 
for obtaining gold from low grade ores, and from the tail- 
ings from other processes. The ore is treated with a 
weak solution of potassium cyanid and allowed to stand, 
exposed to the air, for some hours, when the following 
action takes place: 

2 Au + 4 KCN + O + H 2 = 2 KAu(CN) 2 -f 2 KOH. 

The potassium-gold cyanid thus formed, is dissolved 
out of the ore with water, and decomposed by means of 
zinc, thus: 

2 KAu(CN) 2 + Zn = Zn(CN) 2 , 2 KCN + 2 Au. 

297. Properties and Uses. Pure gold is a bright yel- 
low metal, which is sometimes found crystallized in regular 
octahedrons. It is very soft, and is the most malleable 
and ductile of all the metals. [Compare 223.] It has a 
specific gravity of 19.26, and melts at 1061°. 

Gold is not acted upon by oxygen at any temperature, 
and so remains unaltered in the air. It is not attacked by 
any single acid, (excepting selenic), but dissolves readily 
in aqua-regia forming gold chlorid. 

Gold alloyed with copper or silver, is extensively used 
for coinage, also for making jewelry, and for other orna- 
mental purposes, pure gold being too soft. Most countries , 
except England, use an alloy containing 10 per cent, of 
copper for gold coinage; English gold coin contains 8.34 
per cent, copper. 

The quality of gold alloy is usually expressed in parts 
in twenty-four, these parts being called carats. Pure 
gold is 24 carats fine, while 18-carat gold contains 18 
parts of gold and 6 parts of alloy. For jewelry and other 



238 LECTURE NOTES 

ornamental purposes the quality of the alloy varies, but 
is usually 14-carat. 

Gold is extensively used for gilding various articles. 
This is generally done by direct application of gold leaf, 
although for small articles it is done by electrolysis. 
[Compare 281]. 

298. The Compounds of Gold. Gold forms two series 
of compounds, the aurous, in which the metal is univa- 
lent, and the auric, in which it is trivalent. All these 
compounds are unstable, and when ignited in the air are 
reduced to metallic gold. Most metals, and all reducing 
agents, precipitate gold from a solution of its salts. 

Gold Chlorid, AuCl 3 , is formed when gold is dis- 
solved in aqua-regia. It is a very deliquescent, golden 
yellow, crystalline salt. When hydrochloric acid is added 
to a solution of gold chlorid, and the solution evaporated, 
a compound having the composition HAuCl 4 ,3H 2 0, re- 
mains. This is known as chlor-auric acid, and forms a 
series of well crystallized salts with the alkali-metals. 
Gold chlorid is used in photography for toning silver 
prints. 

An interesting compound of gold, known as the purple 
of Cassius, is formed by treating a dilute solution of gold 
chlorid, with stannous chlorid. Its exact composition is 
not known, but it is believed to contain some fiuely 
divided metallic gold. It has a beautiful purple color, and 
is used in glass and porcelain painting. 

THE METALS OF GROUP II. 

299. The Alkaline Earth Metals. The early chemists 
gave the name earth, to those metallic oxids which were 
light colored, insoluble in water, and did not decompose, 



LECTURE NOTES 239 

or even fuse, at any temperature which they could obtain. 
The oxids of the metals barium, strontium, and calcium, 
closely resembled the earths, except that they were solu- 
ble in water to some extent, and as the solution gave the 
alkaline reaction, they were called alkaline earths. 

The alkaline earth metals constitute division A of Group 
II (276), and were recognized as forming a natural group 
of elements long before the periodic system was discov- 
ered. The metal magnesium has sometimes been consid- 
ered as belonging to this division. Its oxid is very slightly 
soluble in water, and shows a feeble alkaline reaction; 
but in its metallic state, as well as in its salts, magnesium 
much more closely resembles the metals of division B, 
and so is generally regarded as belonging to that division. 

The alkaline earth metals resemble the alkali-metals in 
many ways. They do not occur free in nature; they all 
decompose water at the ordinary temperature, liberating 
hydrogen; their oxids combine directly with water to 
form hydroxids which are strong bases, and which com- 
bine with carbon dioxid to form carbonates; and each 
forms a peroxid. They differ from the alkali-metals in 
that their sulfates, phosphates, and carbonates, are in- 
soluble in water. 

They are all bivalent and form but one class of salts 
with the acids. 

Calcium. Symbol, Ca. Atomic Weight, 40.1. 

300. History and Occurrence. The fact that the an- 
cients used mortar in the construction of their buildings, 
shows that they must have been familiar with the prepa- 
ration and properties of lime, which is one of the most 
important compounds of calcium. The element itself was 
first obtained by Davy, in 1807. 



240 LECTURE NOTES 

Calcium is the fifth most widely distributed element 
known, and constitutes about 3.8 per cent, of the earth's 
crust. It does not occur free in nature, but its com- 
pounds, which are often found in enormous quantities, are 
very numerous, and very widely distributed. Its most 
important natural compound is the carbonate, CaCO s , 
which occurs in the various forms known as calcite, arag- 
onite, chalk, limestone, marble, and coral. Combined 
with magnesium it forms dolomite, or magnesian lime- 
stone, (Ca,Mg)C0 3 . As the sulfate it occurs in the mine- 
rals gypsum, selenite, and anhydrite. Among other im- 
portant minerals may be mentioned fluorite, CaF 2 , and 
apatite, 3 Ca 3 (P0 4 ) 2 , Ca(F,Cl) 2 , and, in addition, calcium is 
found in greater or less quantities in nearly all the sili- 
cates, which make up the bulk of the earth's crust. It is 
also found in the bones and teeth of all animals, in egg 
shells, and in the shells of all mollusks. 

301. Preparation and Properties. Metallic calcium is 
difficult to prepare, and has little or no commercial im- 
portance. It was made by Davy by electrolysis of the 
fused chlorid. It is best prepared by fusing calcium iodid 
with sodium. This forms an alloy of the metals from 
which the sodium is separated by dissolving in pure 
alcohol. 

Calcium is a light yellow metal with a specific gravity 
of 1.5. It is harder than lead, and quite malleable. It 
oxidizes quickly in moist air, and decomposes water rap- 
idly with evolution of hydrogen. When heated to redness 
in the air it burns with a very brilliant light. Many of 
its compounds are of great commercial importance. 

All calcium compounds color the non -luminous gas flame 
a bright orange-red, 



LECTURE NOTES 241 

COMPOUNDS OF CALCIUM. 

302. Calcium Oxid, CaO. This compound, which is 
known commercially as lime, is obtained by strongly ig- 
niting the carbonate. It is made in large quantities for 
commercial purposes, by burning limestone or chalk, 
mixed with some coal, in what is called a lime-kiln. When 
freshly burnt it is called quicklime. 

Pure calcium oxid is a white porous mass. It is infu- 
sible in the oxy-hydrogen flame, but melts in an electric 
furnace, and, on cooling, forms a crystalline mass. When 
lime is heated in the oxy-hydrogen flame it emits a very 
intense white light, which is called the lime-light or Drum- 
mond light. If exposed to the air, lime absorbs moisture, 
increases in bulk, and gradually falls to powder, forming 
calcium hydroxid. If water is poured on quicklime, the 
same process goes on much more rapidly; it increases two 
or three times in bulk, and much heat is evolved. The 
purest lime has the largest increase. This process is 
called slaking and the hydroxid formed is called slaked 
lime. Pure lime when slaked forms a paste with water 
which is unctuous to the touch and so is called fat or 
rich lime. 

Lime is used in the laboratory for drying gases and 
liquids, and in many analytical processes. It is used on 
a very large scale commercially in the preparation of the 
mortars and cements used for building purposes, and for 
the manufacture of bleaching powder. 

303. Calcium Hydroxid, Ca(0H) 2 . This compound, is 
made by the action of water on quicklime. It is a white 
amorphous powder, which dissolves in water with diffi- 
culty, forming what is called lime water. The solution 
has a strong alkaline reaction, and combines readily with 



242 LECTURE NOTES 

carbon dioxid to form calcium carbonate. It is for this 
reason that lime water is used in the laboratory to detect 
the presence of carbon dioxid. 

When slaked lime is mixed with less water than will 
dissolve it, it forms an emulsion called milk of lime. 

304. Mortars and Cements. The products obtained 
from the burning of limestone depend upon the purity of 
the material used, and are known as common or fat 
lime, hydraulic lime, and hydraulic cement. 

Common mortar is made by mixing common slaked lime 
with enough water to make a thin paste, then adding 
three or four parts of sharp sand and thoroughly mixing 
the whole mass. Mortar very soon dries, or "sets," 
enough to give considerable strength to a structure, after 
which, by absorbing carbon dioxid from the air, it gradu- 
ally becomes very hard. The maximum hardness is not 
reached for many years, or even centuries. There is at 
the same time a gradual combination of the lime with the 
silica (sand), forming calcium silicate, which makes the 
mortar more durable. Slaked lime, made into a paste 
with water, can be preserved indefinitely, if kept from 
contact with the air. In making mortar this is commonly 
done, the paste being kept some days before using. 

Hydraulic lime is made from limestone which contains 
some aluminum silicate (clay), often as much as 10 per 
cent. It will slake and fall to powder like common lime, 
but mortar made from hydraulic lime will harden without 
exposure to the air, and will even harden under water. 

Hydraulic cement is made from limestone which usu- 
ally contains more than 20 per cent, of clay. The burning 
of such a limestone results in a more or less fused mass, 



LECTURE NOTES 243 

which does not slake, and so is ground to powder by mill- 
stones. When this powder is mixed with water, it soon 
becomes hard, after which it is not affected by water. 

So few natural limestones contain the right amount of 
clay, that nearly all the cement is made by burning lime- 
stone and clay, mixed together in the proper proportions. 
When this mixture is burned, the lime and clay combine, 
to form calcium silicate, Si0 2 , 3 CaO, and a double silicate 
of aluminum and calcium, Si0 2 (Al 2 3 , CaO) 3 , together with 
some free lime. The "setting" of the cement is believed 
to be due to the combining of these ingredients with the 
water to form a crystalline hydrated silicate of calcium 
and aluminum. 

305. Calcium Fluorid, CaF 2 . This compound occurs in 
nature as the mineral fluorite, often in the form of beau- 
tiful cubical or octahedral crystals. It also occurs in 
small quantities in the bones, and in the enamel of the 
teeth. It is obtained as a voluminous white precipitate 
when a soluble fluorid is added to a solution of calcium 
chlorid. 

Pure calcium fluorid is colorless, but the mineral flourite 
is generally blue, violet, green, or yellow. It is insolu- 
ble in water, and only very strong acids will decompose 
it. If some kinds of fluorite are gently heated, they 
become faintly luminous, long before they approach a red 
heat. This peculiar property is called fluorescence. When 
more strongly heated it becomes p/iospho?*esce?it (142). 
It fuses at a red heat and is used as a flux in the smelting 
of ores. 

306. Calcium Chlorid, CaCl 2 . This compound is found 
in small quantities in sea and river waters, and in the 
mineral tachydrite, . CaCl s ,MgCl 8 , 12 H 2 0, which is found 



244 LECTURE NOTES 

in the Stassfurt salt deposits. It is produced in large 
quantities in many manufacturing processes. (Compare 
125 and 261). 

Calcium chlorid forms large deliquescent crystals hav- 
ing six molecules of water, CaCl 2 , 6H 2 0. When these 
crystals are heated to 200° they lose all their water, and 
form an extremely hygroscopic, porous mass, which is 
much used in the laboratory for drying gases and liquids. 
The crystallized salt is very soluble in water, and the 
solution is attended with considerable reduction of tem- 
perature. When it is mixed with ice or snow the 
temperature may be lowered to — 48°. 

Calcium chlorid combines directly with dry ammonia 
forming the compound CaCl 2 ,8NH 3 , and so it cannot be 
employed for drying ammonia gas. 

307. Bleaching Powder, Ca(C10)Cl. This important 
compound, which is known commercially as chlorid of 
lime, is manufactured in enormous quantities by the action 
of chlorin on slaked lime, thus: 

2 Ca(OH) 2 -f 2 CI, = CaCl 2 + Ca(C10) 2 + 2 H 2 0. 

It is produced in large quantities as a by-product in 
theLeblanc soda process, and forms an important addition 
to that industry. The hydrochloric acid, obtained in the 
salt-cake process, is decomposed by manganese dioxid 
giving chlorin (64), which is used for making the bleach- 
ing powder. Bleaching powder is also made from the 
chlorin obtained by Castner's process for making sodium 
hydroxid (244). 

In the manufacture of the bleaching powder, the slaked 
lime is spread upon the floor of a large chamber, and ex- 
posed to the action of the chlorin. The action takes place 



LECTURE NOTES 245 

as expressed by the equation given above, and is usually 
complete in about two days. 

It was at first believed that bleaching powder was a 
mixture of calcium hypochlorite, Ca(C10) 2 , and calcium 
chlorid, CaCl 2 , but it has been shown that the substance 
does not contain free calcium chlorid. It is at present 
believed to be a chemical combination of the two salts 
which may be expressed by the formula, 

n /CI 
0a <O-Cl. 

It is a white porous solid, with an odor like that of 
chlorin. It is quite soluble in water, and the solution has 
an alkaline reaction. Carbon dioxid decomposes it, libe- 
rating hypochlorous acid, so that it cannot be preserved 
in contact with the air. It decomposes slowly, even in 
closed vessels, with liberation of oxygen. Dilute acids 
decompose it with liberation of chlorin. 

The commercial value of bleaching powder depends upon 
the amount of available chlorin which it contains. Good 
bleaching powder should contain at least 25 per cent., and 
the amount sometimes reaches as high as 38 per cent. 
Its principal uses are for bleaching and disinfecting. 
(Compare 73). 

308. Calcium Sulfate, CaS0 4 . This compound occurs 
in nature as the mineral anhydrite, CaS0 4 , and gypsum, 
CaS0 4 , 2 H 2 0, of which there are several varieties known 
as selenite, alabaster and satin spar. It is prepared by 
adding sulfuric acid or a soluble sulfate, to a solution of 
calcium chlorid. 

Prepared calcium sulfate is a white crystalline powder, 
which is slightly soluble in water. It is used as a filling 
in the manufacture of writing paper. 



246 LECTURE NOTES 

When gypsum is heated to about 120°, it loses a part 
of its water of crystallization, and is converted into what 
is called burnt gypsum, or plaster of Paris, 2 CaS0 4 ,H 2 0. 
When this substance is mixed with water, it evolves heat, 
and gradually becomes hard. The hardening of the plaster, 
as it is commonly called, is due to the taking on of water 
of crystallization, and so re-forming gypsum. 

Plaster of Paris is largely used for ornamental plaster 
work, for making plaster casts, and as a cement. 

309. Calcium Phosphate, Ca 3 (P0 4 ) 2 . This compound 
occurs in nature in a number of minerals, containing 
more or less impurity, of which the most important are 
apatite, phosphorite, and sombrerite. It is the chief con- 
stituent of the phosphate beds, which are found as im- 
mense deposits in various parts of the earth. It is the 
chief inorganic constituent of bones, burnt bones con- 
taining about 80 per cent, of this substance. 

Calcium phosphate is insoluble in water, but dissolves 
in water containing in solution certain salts, such as 
sodium chlorid or nitrate. It is also soluble to some ex- 
tent in all acids, even in carbonic acid. 

It is a valuable fertilizer and is extensively used for 
this purpose. For the reasons just given it dissolves 
easily in the soils, and is taken up by the plants, where it 
is found in the seeds and fruit. From these it finds its 
way into the animal body. 

310. Calcium Carbid, CaC 2 . This compound has been 
knowD for a long time, but has only recently become a 
commercial product. It is now made on a large scale by 
heating a mixture of lime and coke in an electric furnace, 
thus: 

CaO + 3 C = CaC 3 -f CO. 



LECTURE NOTES 247 

Calcium carbid is a hard, grayish-black, crystalline 
substance. It is easily decomposed by water, forming 
acetylene (187), which is much used for illuminating 
purposes. 

311. Calcium Carbonate, CaC0 3 . This is undoubtedly 
the most important compound of calcium, since it is not 
only of great commercial value in the forms in which it is 
found in nature, but it is also used, directly or indirectly, 
in the preparation of all the compounds of calcium. 

In the various forms of limestone and marble, it is found 
in nature in enormous quantities, often comprising whole 
mountain ranges. It forms the greater part of the shells 
of all mollusks. Chalk, which is found in large deposits, 
is composed chiefly of the shells of microscopic marine 
animals, principally foraminifera, and is nearly pure cal- 
cium carbonate. The shells of the coral animal, which 
form the coral islands, are also calcium carbonate. 

Calcium carbonate is practically insoluble in pure 
water, but dissolves to a considerable extent in water 
containing carbon dioxid, forming in this way the acid 
calcium carbonate, which is so commonly found in all 
natural waters (Compare 59). It is decomposed by 
nearly all acids liberating carbon dioxid. 

The natural limestones and marble, are used very ex- 
tensively for building stone, and for ornamental purposes, 
and in the manufacture of most other calcium compounds. 

Strontium. Symbol, Sr. Atomic Weight, 87.6. 

312. Occurrence, Preparation, and Properties. This 
metal, which is quite rare, is found principally in the 
minerals strontianite, SrC0 3 , and celestite, SrSO^. 

The metal was first prepared by Davy in 1807 by elec- 



248 LECTURE NOTES 

trolysis of the chlorid. It is best made by a method 
analogous to that used for the preparation of calcium. 

It is a light yellow metal, having a specific gravity of 
2.5. It is harder than lead and quite malleable. It oxi- 
dizes quickly in moist air, and decomposes water violently, 
even in the cold, with evolution of hydrogen. 

313. The Compounds of Strontium. The compounds 
of strontium are almost exactly analogous to those of cal- 
cium, and may be formed in the same way. Owing to 
their rarity they are not very much used. They are 
mostly obtained from strontianite. 

Strontium nitrate, Sr(N0 3 ) 2 , is made by dissolving the 
carbonate in nitric acid. When strontium nitrate is 
heated with carbon, the mixture burns with a character- 
istic, carmine-red flame. It is therefore largely used in 
pyrotechny for the production of red fire. 

Barium. Symbol, Ba. Atomic Weight, 137.4. 

314. Occurrence, Preparation, and Properties. Barium 
is found in nature in large masses in the minerals barite, 
BaS0 4 and witherite, BaC0 3 , and, in small quantities, in a 
number of other minerals. These, as well as all other 
barium compounds, are distinguished by their high speci- 
fic gravity. 

Metallic barium is very difficult to prepare, and it is a 
question if it has ever been obtained pure. It was first 
obtained by Davy by electrolysis. It is best prepared by 
adding- sodium amalgam to a hot saturated solution of 
barium chlorid, and distilling off the mercury, from the 
barium amalgam thus formed, in a stream of hydrogen. 

Barium is a yellow (?) metal, having a specific gravity 
of 3.6. It fuses with great difficulty, but decomposes 



LECTURE NOTES 249 

water with great energy at the ordinary temperature, 
liberating hydrogen. It is the most basic of any of the 
metals of this group. 

COMPOUNDS OF BARIUM. 

315. The Oxids and Hydroxid of Barium. Barium 
forms two oxids, and one hydroxid, which are exactly 
analogous to the corresponding compounds of calcium. 

Barium oxid, BaO, which is also called baryta, is best 
prepared by heating barium nitrate. When heated to a 
low red heat, in a stream of air or oxygen, it combines 
with the latter forming barium dioxid. 

Barium hydroxid, Ba(OH) 2 , is formed when the oxid is 
dissolved in water. It is best formed by heating the car- 
bonate, together with a little carbon, in a current of 
steam. It is quite soluble in hot water, the solution be- 
ing called baryta-water. It is more strongly alkaline 
than lime-water, and is used in the laboratory for the 
same purposes. 

Barium dioxid, or peroxid, Ba0 2 , is made by heating- 
barium oxid in the air or oxygen. It is a gray powder, 
which at a strong red heat gives up oxygen and forms the 
monoxid. It is thus employed in Brin's process for ob- 
taining oxygen from the air. It is also used for the 
preparation of hydrogen dioxid (60). 

316. Barium Chlorid, BaCl 2 . This compound is made 
by dissolving the mineral witherite in hydrochloric acid. 
It is also obtained on a large scale from the mineral ba- 
rite, by heating a mixture of barite, carbon, calcium 
chlorid and limestone in a reverberatory furnace, thus: 

BaS0 4 + 4 C + CaCl 2 = BaCl 2 + CaS + 4 CO. 



250 LECTURE NOTES 

The mass is lixiviated with water, when barium chlorid 
dissolves, and insoluble calcium sulfid remains behind. 

Barium chlorid forms colorless tabular crystals with 
two molecules of water, BaCl 2 , 2 H 2 0. It has a bitter 
taste and is quite poisonous, as are all the soluble barium 
salts. It is the principal source of barium compounds in 
the laboratory. 

317. Barium Sulfate, BaS0 4 . This compound occurs 
in large quantities in nature as the mineral barite. It is 
formed as a heavy white precipitate whenever sulfuric 
acid, or any soluble sulfate, is added to a solution of a 
barium salt. 

Barium sulfate is a very heavy compound, its specific 
gravity being about 4.5. It is insoluble in water, and 
practically so in all other solvents, except hot concen- 
trated sulfuric acid, in which it dissolves to a slight ex- 
tent. Precipitated barium sulfate is largely used as a 
pigment. It is manufactured on a large scale and known 
as permanent white. It is used as a substitute for white 
lead. 

318. Barium Nitrate, Ba(N0 3 ) 2 . This compound is 
obtained by dissolving the mineral witherite in nitric 
acid. It is also formed by mixing hot concentrated solu- 
tions of sodium nitrate and barium chlorid, thus: 

NaNO s + CaCl 2 = 2 NaCl + Ba(N0 3 ) 2 . 

On cooling the mixture, barium nitrate crystallizes out 
in colorless octahedrons. 

Its principal use is in pyrotechny, in which it is used 
for the preparation of green fire. 

319. Barium Carbonate, BaC0 3 . This compound occurs 



LECTURE NOTES 251 

in nature as the mineral witherite. It can be prepared by- 
precipitating a solution of a barium salt with a soluble 
carbonate. 

Prepared in this way it is a white amorphous powder, 
which is used in chemical analysis. The natural compound 
is used in the preparation of other barium compounds. 

320. The Magnesium Group of Metals. The remain- 
ing metals of Group II, which belong to division B. are 
not quite so closely related to each other as are the alka- 
line earth metals. These metals are giucinum, magnesium, 
zinc, cadmium, and mercury. They are all bivalent, and, 
in addition, mercury also forms a series of salts in which 
it is univalent. In its compounds giucinum much resem- 
bles aluminum, while the univalent salts of mercury are 
much like the corresponding salts of silver and copper. 
Mercury is also peculiar in being a liquid at the ordinary 
temperature. The metals are all quite easily obtained, 
and at the ordinary temperature are permanent in the 
air. None of them will decompose water at the ordinary 
temperature but magnesium and zinc will do so at a high 
temperature. When strongly ignited they will all combine 
with oxygen to form oxids, although the oxid of mercury 
is decomposed at a still higher temperature, liberating 
oxygen (47). The oxids and hydroxids are insoluble in 
water while the sulfates are all soluble. 

Giucinum. Symbol, Gl. Atomic Weight, 9. 

321. Giucinum and its Compounds. This rare element, 
which is more commonly called beryllium, occurs in nature 
in a number of minerals of which the most important are 
beryl, 3 BeO, Al 2 O s , 6 Si0 2 , and chrysoberyl, BeO, Al 2 O s . 
When beryl is transparent and dark-green, or bluish- 
green, it is used as a gem under the names emerald and 



252 LECTURE NOTES 

aquamarine, respectively. The metal was first obtained 
by Wohler in 1828 by fusing the chlorid with potassium. 
It is best obtained by passing the vapor of the chlorid 
over melted sodium in a glass tube, all the air having 
been replaced by hydrogen. 

Glucinum is a silver- white metal, which much resembles 
magnesium. It has a specific gravity of 2.1. It remains 
unchanged in the air at the ordinary temperature, but 
when strongly heated, it becomes covered with a film of 
oxid; and if the powdered metal is heated, it takes fire 
and burns with great brilliancy. It is soluble in the 
common acids, especially when warm, and in sodium and 
potassium hydroxids. It does not decompose water even 
at a red heat. 

The compounds of glucinum most resemble those of 
aluminum. They have a peculiar sweetish taste, hence 
the name glucinum. 

Magnesium. Symbol, Mg. Atomic Weight, 24.3. 

322. History and Occurrence. Compounds of this metal 
have been known since the early part of the seventeenth 
century. Metallic magnesium was first prepared by Davy, 
although he did not succeed in obtaining it pure. 

Magnesium is the sixth most widely distributed element 
known, and constitutes about 2.7 per cent, of the earth's 
crust. It does not occur free in nature, but its com- 
pounds are quite numerous, and very abundant. Its most 
abundant compound is the mineral dolomite, (Mg,Ca)CO s , 
which occurs in mountainous masses. Some of the 
other important minerals which contain it are magnesite, 
MgC0 3 , kieserite,MgSO v Hfl, carnallite,MgC\ 2 ,KC\,6 H 2 0, 
and kainite, MgS0 4 , KC1, 6 H 2 0. It is an important con- 
stituent of many of the silicates, such as the pyroxenes, 



LECTURE NOTES 253 

hornblendes, micas, serpentine, and talc, while in smaller 
quantities it is found in most of the other silicates. Its 
soluble salts are found in sea water, and in almost all 
natural waters. 

323. Preparation and Properties. Magnesium can be 
prepared by the electrolysis of the fused chlorid, or by 
heating the chlorid with sodium. It is best prepared by 
strongly heating in a closed crucible, a fused mixture of 
magnesium and sodium chlorids, with metallic sodium, 
together with some powdered fluorite to act as a flux, thus: 

MgCl 2 + NaCl + 2 Na = 3NaCl -f-Mg. 

Magnesium is a brilliant, silver-white metal, which is 
quite malleable and ductile. It has a specific gravity of 
1.75, melts at about 700°, and boils at about 1100°. It 
oxidizes slowly in moist air and slowly decomposes boiling 
water. When magnesium is heated to redness in the air, 
it takes fire and burns with a white light of dazzling bril- 
liancy. This light is very rich in the chemically active rays, 
and, for this reason, is employed in photography for 
artificial illumination. Magnesium combines directly 
with nitrogen, when strongly heated in that gas, and 
forms magnesium nitrid, Mg 3 N 2 (Compare 117). 

Magnesium in the form of powder is a powerful reduc- 
ing agent, and reduces most metallic oxids when heated 
with them (Compare 206 and 216). The intense light of 
burning magnesium is sometimes used in signalling, and 
also in pyrotechny. 

COMPOUNDS OF MAGNESIUM. 

324. Magnesium Oxid, MgO. This compound, which is 
also called magnesia, occurs in nature as the mineral peri- 
clase. It is formed when magnesium burns in the air. 



254 LECTURE NOTES 

It is also formed on a largo scale by heating the carbonate. 
It is a white, very voluminous, amorphous powder, and is 
known commercially as magnesia usta, or calcified magnesia. 

When magnesia is heated, it becomes highly luminous, 
and also conducts electricity. If. therefore, a cylinder of 
magnesia is heated to the point of conductivity, it can be 
used for electric lighting. This is the principle of the 
new Nernst incandescent electric lamps. Magnesia is an 
extremely refractory substance, and so is used for mak- 
ing firebricks, and crucibles, and for lining furnaces. 

325. Magnesium Hydroxid, MgvOH\. This compound 
occurs in nature as the mineral brucite. It is obtained as 
a white gelatinous precipitate, when sodium or potassium 
hydroxid is added to a solution of a magnesium salt. It 
is very nearly insoluble in water, but easily dissolves in 
water containing ammonium salts in solution. It slowly 
absorbs carbon dioxid from the air. forming a carbonate. 
and when heated forms the oxid. If the heat employed 
is not too strong, the oxid formed can be used as a cement, 
since on mixing with water, it will rehydrate itself, and 
harden in a few hours. 

326. Magnesium Chlorid, MgCl.,. This compound is 
found in considerable quantities in sea water and many 
mineral springs. It occurs combined with potassium 
chlorid. as the mineral carnallite in the Stassfurt salt de- 
posits. It is produced in large quantities in various 
technical processes. It can be prepared by dissolving the 
oxid. hydroxid or carbonate in hydrochloric acid. 

Magnesium chlorid is very soluble in water and when 
this solution is evaporated the salt crystallizes out in pris- 
matic crystals, with six molecules of water. MgCl 3 , 6 H s O. 



LECTURE NOTES 255 

It forms a series of double salts with the alkaline chlorids, 
which have the general formula MgCl 2 ,M'Cl, 6 H a O. 

327. Magnesium Sulfate, MgS0 4 . This important com- 
pound was discovered in 1618 in the water from a mineral 
spring at Epsom, in England, and so was called Epsom 
salts, a name by which it is still known commercially. 
It is found in sea water and many mineral springs, and 
as the mineral kieserite, K 2 S0 4 , H 2 0, in the salt deposits 
at Stassfurt, from which it is largely obtained. It is also 
obtained by decomposing dolomite, (Mg,Ca)C0 3 , with sul- 
furic acid. The decomposed mass is lixiviated with water 
when the soluble magnesium sulfate dissolves, leaving the 
insoluble calcium sulfate. 

Magnesium sulfate crystallizes from an aqueous solution 
with seven molecules of water, MpSO,, 7 HO. If these 
crystals are heated, they lose six molecules of water below 
150°, while the seveuth does not pass off until 200°. It 
is believed that this last molecule of water, being much 
more difficult to remove, exists in the compound as water 
of constitution, the following constitutional formula being- 
given for this compound, viz. : 

H—O— Mg— (X q // 

A number of other crystalline salts containing water 
show this same peculiarity and may be explained in the 
same way. This is believed to be the constitutional for- 
mula of the mineral kieserite, and may explain the fact, 
that kieserite is much less soluble in water than the com- 
mon crystalline form. 

Magnesium sulfate forms transparent prismatic crys- 
tals, which have a bitter taste and are much used in 



256 LECTURE NOTES 

medicine as a safe cathartic. It forms a series of double 
salts with the alkaline sulfates. 

328. The Monoclinic Double Salts. The sulfates of the 
metals magnesium, zinc, and cadmium, which belong to 
this group, together with copper sulfate, and those sul- 
fates of manganese, iron, nickel and cobalt in which the 
metals are bivalent, each form with the alkaline sulfates 
a series of double salts. These crystallize with six mole- 
cules of water, and are isomorphous, that is, they have 
the same crystalline form, with the same crystalline con- 
stants, and are capable of overgrowth. The crystalline 
form is a modified prism belonging to the monoclinic 
system of crystallography, and so these crystals are 
called the monoclinic double salts. 

They have the general formula M"S0 4 ,M/S0 4 , 6 H 2 0, 
where M" is one of the bivalent metals given above, and 
M' is some univalent metal, usually an alkali-metal or 
ammonium. The following is believed to be their general 
constitutional formula, viz. : 

M '~ °>S X/ ° 

W' ^ ^ 6 H O 

M'— 0^ b ^ O 

329. Magnesium Carbonate, MgC0 3 . This compound 
occurs, in nature as the mineral magnesite. If sodium or 
potassium carbonate, is added to a solution of a magne- 
sium salt, a white basic carbonate is formed. If this 
compound is carefully dried, a voluminous white powder re- 
main s which has the composition 3 MgC0 3 , Mg(OH) 2 , 4 H 2 0. 
This compound is much employed both commercially and in 
medicine under the name magnesia alba. If magnesia 



LECTURE NOTES 257 

alba is suspended in water and carbon dioxid passed 
through the liquid, the normal carbonate can be formed. 
Magnesium carbonate, both normal and basic, is insoluble 
in water, but dissolves in water containing ammonia. 
If either carbonate is strongly heated it decomposes, 
forming magnesium oxid and carbon dioxid. 

Zinc. Symbol, Zn. Atomic Weight, 65.4. 

330. History and Occurrence. The alloy of copper and 
zinc which we call brass, was known to the ancients, 
although they did not recognize it as an alloy, but sup- 
posed it to be a peculiar kind of copper. The alchemists 
knew that when copper was heated with certain ores, now 
known to contain zinc, it became yellow. They did not 
distinguish between zinc and tin, nor between zinc and 
bismuth, and it was not until the latter part of the seven- 
teenth, or early part of the eighteenth century, that zinc 
was recognized as a distinct metal. 

Zinc is said to have been found native in small quanti- 
ties near Melbourne, Australia. With this exception it 
is always found in combination, the most important ores 
being smithsonite, ZnCO s , calamine, (ZnOH) 2 Si0 3 , and 
sphalerite, or zinc-blende, ZnS. There are, in addition to 
these, three minerals found in New Jersey in such quan- 
tities as to make them important zinc ores. These are 
zincite, ZnO, franklinite, (Zn,Fe)0,F 2 O s , and willemite, 
Zn 2 Si0 4 . 

331. Preparation. In the preparation of zinc, the ores, 
other than the oxid, are first roasted and thus changed to 
the oxid. The oxid is then mixed with coal and heated 
to a high temperature in earthenware retorts, forming the 
metal, thus: 

ZnO -j- C = Zn -f CO. 



258 LECTURE NOTES 

On further heating the zinc distils over and is collected 
in iron receivers. 

There are two processes employed for the reduction of 
zinc ores, which differ only in unimportant details, the 
chemical action being the same in each. In the Silesian 
process, which is the one more generally employed, the 
distillation is carried on in large retorts, while in the 
Belgian process, long tubes are used for the distillation. 

During the process of distillation, some of the zinc be- 
comes partially oxidized and passes over into the receiver 
in the form of a fine dust. This is a mixture of zinc oxid 
and finely divided metallic zinc, and is called zinc-dust. 
This is employed as a reducing agent in organic chemis- 
try, and in certain technical processes. 

332. Properties and Uses. Pure zinc is a bluish- white, 

crystalline metal. At the ordinary temperature it is quite 
brittle, but between 100° and 150° it may be rolled into 
thin sheets, and drawn into wire. At a little above 200° 
it becomes very brittle, and may be powdered in a mortar. 

Zinc has a specific gravity of 6.9. It melts at 420° and 
boils at about 940°. The density of its vapor at about 
1200° shows that its molecule is monatomic, that is, con- 
tains but a single atom (Compare 14 and 19). 

Zinc is permanent in dry air at ordinary temperatures, 
but becomes slightly tarnished on the surface in moist air. 
If heated in the air to a little above its melting point, it 
takes fire and burns with an intense bluish- white flame. 

Pure zinc is not easily acted upon by acids, but the 
presence of even small quantities of any impurity causes it 
to dissolve rapidly, so that ordinary commercial zinc dissolv- 
es easily in dilute acids, with evolution of hydrogen. It is 



LECTURE NOTES 259 

also soluble in a hot solution of sodium or potassium hy- 
droxid with evolution of hydrogen, thus: 

Zn + 2 NaOH = Zn(ONa), + H 2 . 

Commercial zinc will decompose water at the boiling 
point, while the pure metal will only do so at a red heat. 

Zinc is extensively used for "galvanizing" iron, which 
is done by dipping the iron into molten zinc, and not by 
electric deposition, as the name might imply. The layer 
of zinc keeps the iron from rusting, and hence galvanized 
iron is used for making articles which are especially ex- 
posed to the action of air and water. It is also used very 
extensively in the manufacture of the important alloy 
called brass (229), which is composed of copper and zinc, 
and sometimes contains a small amount of tin or lead. 

COMPOUNDS OF ZINC. 

333. Zinc Oxid, ZnO. This compound occurs in nature 
as the mineral zincite. It is formed when zinc burns in 
the air, and is manufactured on a large scale in this way 
for commercial purposes. 

Zinc oxid is a soft white substance which was known to 
the alchemists under the name philosopher's wool. It is 
known commercially as zinc white, and is very largely 
used as a pigment in place of the more common "white 
lead." Paint made from zinc white is superior to lead 
paint for some kinds of work, since it is not blackened by 
hydrogen sulfid. 

It is infusible in the oxy-hydrogen flame, but after be- 
iug heated remains phosphorescent for some time. It 
acquires a yellow color when hot. but loses this and 
becomes white again on cooling. 



260 LECTURE NOTES 

334. Zinc Chlorid, ZnCl 2 . This compound is formed 
when zinc, or its oxid, or carbonate, is dissolved in hydro- 
chloric acid. 

Zinc chlorid is a soft white substance which melts at a 
little above 100°, and boils at 730°. It is very deliquescent, 
and is easily soluble in water and in alcohol. It is used in 
surgery as a caustic. When mixed with zinc oxid, and 
made into a paste with water, it quickly sets to a hard 
mass. It is used for the preservation of wood, and as a 
disinfectant. 

335. Zinc Sulfate, ZnS0 4 . This compound occurs in a 
few localities as the rare mineral goslarite. It is formed 
in the passages of certain zinc mines by the oxidation of 
sphalerite (ZnS). It was first prepared by roasting zinc 
ores containing sphalerite, lixiviating the roasted mass 
with water, and separating the zinc sulfate by crystalli- 
zation. It was known to the alchemists as ivhite vitriol 
(106). 

Zinc sulfate forms white prismatic crystals, which con- 
tain seven molecules of water, ZnS0 4 , 7 H 2 0. They efflo- 
resce slowly in the air, and lose six molecules of water at 
100°, but the seventh only at a red heat (Compare 327). 
The crystals are isomorphous with the corresponding ones 
of magnesium sulfate. 

Zinc sulfate has an acrid, metallic taste. It forms a 
series of monoclinic double salts with the alkaline sul- 
fates (328). It is used in medicine and in dyeing. 

336. Other Compounds of Zinc. Some of the other 
compounds of zinc are found in nature, and are used as 
ores. The rest are not of especial importance. 

Zinc sulfide Zn.S, occurs as the mineral sphalerite, and 



LECTURE NOTES 261 

is formed as a white precipitate, when a solution of a zinc 
salt is treated with a soluble sulfid. 

Normal zinc carbonate, ZnCO t , is found in nature as 
the mineral smithsonite. When a solution of a zinc salt is 
treated with a soluble carbonate, a basic zinc carbonate, 
usually (ZnOH) 2 C0 3 , is formed. 

Cadmium. Symbol, Cd. Atomic Weight, 112.4 

337. History and Occurrence. Cadmium is a somewhat 
rare metal. It was discovered in a zinc ore by Stromeyer 
in 1817. 

Cadmium does not occur free in nature, and the only nat- 
ural compound of which it is the chief constituent, is the very 
rare mineral greenockite, CdS. It occurs in small quanti- 
ties in many zinc ores especially in sphalerite and smithson- 
ite and these ores are the principal sources of cadmium. 
The name cadmium comes from its association with zinc, 
zinc oxid having been called cadmia by the alchemists. 

338. Preparation, Properties and Uses. In the process 
of extracting the zinc, from its ores which contain cad- 
mium, the latter being more volatile, and more easily 
oxidized, is found in the first portions of the distillate, 
partly as metal, and partly as oxid. This portion is dis- 
solved in hydrochloric or sulfuric acid, and the cadmium 
precipitated with metallic zinc, thus: 

CdS0 4 + Zn = ZnS0 4 + Cd. 

Cadmium is a white, tenacious, rather soft metal, some- 
what resembling zinc, but much more malleable and 
ductile. It has a specific gravity of 8.6. It melts at 
320°, and boils at 770°. Determination of its vapor den- 
sity shows that its molecule is monatomic, like that of zinc. 



262 LECTURE NOTES 

When cadmium is heated in the air, it takes fire and 
burns with a bright light, forming brown fumes of cad- 
mium oxid. It dissolves slowly in hydrochloric or sulfuric 
acid, but easily in nitric acid. 

Cadmium is chiefly employed in making the so-called 
fusible alloys, such as Wood's metal (See 228). An amal- 
gam of cadmium hardens quickly, and was formerly used 
for filling teeth. 

339. The Compounds of Cadmium. The compounds of 
cadmium are quite similar in their properties to the cor- 
responding compounds of zinc. On account of their 
comparative rarity they are not much used in the arts. 
They can all be obtained by the usual methods. 

Cadmium iodid, Cdl 2 , can also be obtained by the direct 
union of cadmium and iodin in the presence of water. It 
forms transparent, tabular crystals, which are soluble in 
water, and in alcohol. It is used in photography, it being 
one of the few iodids which are soluble in alcohol. 

Cadmium sulfid, CdS, occurs in nature as the mineral 
greenockite. It is a yellow compound and employed as a 
pigment. 

Cadmium sulfate, CdS0 4 , crystalizes from an aqueous 
solution at the ordinary temperature, in crystals which 
have the composition 3 CdS0 4 ,8H 2 0. It forms a series of 
monoclinic double salts with the alkaline sulfates (Com- 
pare 328). 

Mercury. Symbol, Hg. (Hydrargyrum). Atomic Weight, 

200. 

340. History and Occurrence. Mercury is one of the 
metals which have been known since very early times. It 
was first mentioned by Theophrastus, about B. C. 300, 



LECTURE NOTES 263 

who referred to it as liquid silver or quicksilver. It de- 
rived its name from the god Mercury. The alchemists, 
who associated the metals known to them with the planets, 
naturally associated this metal with the planet Mercury, 
and gave it the symbol Q . This symbol is supposed to 
have been intended to represent the "caduceus," or official 
wand, of the god Mercury. The alchemists knew how to 
purify the metal by distillation, and how to prepare some 
of its compounds. 

Mercury is found in the native state, in the form of 
small globules scattered through its ores. It is also found 
occasionally as an amalgam with gold and silver. Its 
principal ore is the mineral cinnabar, HgS, the most im- 
portant mines of which are in Almaden, in Spain ; in Idria, 
in Austria; in New Almaden and New Idria in California; 
and in China and Japan. 

341. Preparation. The process of extracting mercury 
from its ores is very simple. The ore is roasted in a 
reverberatory furnace, where the sulfur is oxidized and 
passes off as sulfur dioxid, while the mercury, which is 
first volatilized, is condensed in cold chambers. The 
action is shown by the following equation: 

HgS + O, = SO, + Hg. 

Another method which is sometimes used, is to mix 
the ore with lime and heat in a closed retort, when the 
mercury distils over, thus: 

4 HgS + 4 CaO = 3 CaS + CaS0 4 -f 4 Hg. 

The mercury thus obtained, usually contains a small 
amount of other metals in solution, such as zinc, tin, lead, 



264 LECTURE NOTES 

and others, and may be purified by distillation, either in 
the air or in a vacuum. For most purposes it may be 
obtained sufficiently pure by allowing a thin stream of the 
impure metal to fall into along tube filled with nitric acid, 
which has been diluted with twenty volumes of water. 
This operation may be repeated until the metal is pure, 
after which it should be washed and dried. 

342. Properties and Uses. Mercury is a silver-white 
metal, and has a brilliant metallic luster. It has a specific 
gravity of 13.59 at 0°. It melts at — 39°, and boils at 
357°. The density of its vapor shows that its molecule is 
monatomic, like that of zinc and cadmium. 

Mercury is the only metal which is a liquid at the ordi- 
nary temperature; when frozen, it is malleable, and crys- 
tallizes in octahedrons. It unites with most metals to 
form alloys which are called amalgams (See 230). 

Mercury remains unchanged when exposed to the air at 
ordinary temperatures, but near the boiling point it grad- 
ually becomes oxidized, forming mercuric oxid. It is not 
attacked by hydrochloric acid, but dissolves in hot con- 
centrated sulfuric acid with evolution of sulfur dioxid. It 
dissolves easily in strong nitric acid forming mercuric 
nitrate, and slowly in dilute nitric acid forming mercurous 
nitrate. The vapor of mercury is very poisonous. 

Mercury is a very useful metal, being employed in 
metallurgical processes (Compare 289 and 296), and in 
the electrolytic process for making sodium hydroxid (Com- 
pare 244). Since it is a liquid at all ordinary tempera- 
tures, it is also used in the manufacture of certain physical 
and chemical apparatus (Compare 11 and 12), and in many 
other ways. 



LECTURE NOTES 265 

COMPOUNDS OF MERCURY. 

343. The Mercurous and Mercuric Compounds. Mer- 
cury forms two series of compounds which are known as the 
mercurous and mercuric compounds. The first are analo- 
gous to those of silver and univalent copper, the metal 
being univalent, and the second are like the compounds of 
zinc and cadmium, the metal being bivalent. When a 
substance reacting with mercury is in excess, mercuric 
compounds are usually formed; and when the mercury is 
in excess, mercurous compounds are the result. Oxidiz- 
ing agents change mercurous compounds to mercuric, 
while reducing agents change mercuric compounds to 
mercurous. 

The mercurous compounds are, as a rule, much less 
stable than the corresponding mercuric compounds. All 
the compounds of mercury are poisonous, the mercuric 
compounds being more poisonous than the corresponding 
mercurous compounds. 

THE MERCUROUS COMPOUNDS. 

344. Mercurous Chlorid, HgCl. This compound which 
is commonly called calomel, is found in small quantities 
in nature. It is commonly made by heating mercuric 
chlorid with mercury, when mercurous chlorid sublimes as 
a white mass of radiating crystals, thus: 

HgCl, + Hg = 2 HgCl. 

It is always formed as an amorphous white precipitate 
when hydrochloric acid, or a soluble chlorid, is added to 
a solution of a mercurous salt. 

Mercurous chlorid is insoluble in water, and in dilute 
acids, but if it is boiled with hydrochloric acid, or certain 



266 LECTURE NOTES 

chlorids, it decomposes, forming mercuric chlorid and free 
mercury. 

When mercurous chlorid is heated, it sublimes without 
fusing. It is a heavy compound, its specific gravity being 
nearly 7. When ammonium hydroxid is added to mercu- 
rous chlorid, it forms a black compound having the 
composition Hg NH CI, and called !amido-mercurous chlo- 
rid. From this action is derived the name calomel. 

345. Mercurous Nitrate, HgNO s . This compound is 
produced when an excess of mercury is acted upon by cold 
dilute nitric acid. It forms large monoclinic crystals 
containing two molecules of water of crystallization, and 
is the principal source of mercurous compounds in the 
laboratory. It dissolves in water containing nitric acid, 
but pure water partially decomposes it and forms a yellow 
basic salt having the composition, HgN0 3 , HgOH. If 
either the normal or basic nitrate is boiled with an excess 
of mercury, it forms a white basic salt, having the com- 
position 3 HgNO g , 2 HgOH. On boiling with water this 
latter salt decomposes into mercury and mercuric nitrate. 

The nitric acid solution of mercurous nitrate gradually 
changes to mercuric nitrate. This may be prevented by 
adding metallic mercury to the solution, when the mercu- 
rous nitrate is again formed. 

The other mercurous compounds are generally unstable 
and of little importance. 

THE MERCURIC COMPOUNDS. 

346. Mercuric Oxid, HgO. This compound was known 
to the alchemists as red precipitate, and is commonly 
known as red oxid of mercury. It may be prepared by 
heating mercury in the air for a long time nearly to the boil- 



LECTURE NOTES 267 

ing point. It is best prepared by heating a mixture of 
mercuric nitrate and mercury. 

When prepared by either of these methods, mercuric 
oxid is a red crystalline powder. When sodium or potas- 
sium hydroxid is added to a solution of a mercuric salt, 
mercuric oxid is obtained as an orange-yellow, amorphous 
precipitate, which contains water, and is believed by some 
to be mercuric hydroxid, Hg(OH) 2 . On being heated, it 
is changed to the red variety. 

When mercuric oxid is heated, it becomes darker, and 
gradually becomes almost black, but on cooling, it assumes 
its original color. At a red heat it decomposes com- 
pletely into mercury and oxygen (47). 

347. Mercuric Chlorid, HgCl 2 . This compound was 
known to the early alchemists, and was called corrosive 
sublimate, a name by which it is still known commercially. 
It can be formed by dissolving mercury in aqua regia, or 
by dissolving the oxid in hydrochloric acid. 

It is prepared commercially by heating a mixture of 
mercuric sulfate and sodium chlorid, (together with a little 
manganese dioxid to prevent the formation of mercurous 
chlorid), thus: 

HgS0 4 + 2 NaCl = HgCl 2 -f Na,S0 4 . 

Mercuric chlorid is a white substance, which dissolves 
easily in water, from which solution it crystallizes in 
rhombic needles. It has a specific gravity of 5.4. It 
melts at 288°, and boils at 303°. It is a violent poison, 
the best antidote being albumen, with which it forms an 
insoluble - compound. It is one of the best antiseptics 
known, and for this reason is much used by taxidermists. 



268 LECTURE NOTES 

When ammonium hydroxid is added to a solution of 
mercuric chlorid, it forms a white compound which has 
the composition HgNH 2 Cl, and is called amido-mercuric 
chlorid (Compare 127). This compound was known to the 
alchemists, and called by them white precipitate. 

Mercuric chlorid forms a number of double salts with 
the alkaline chlorids, which are crystallizable compounds, 
and easily soluble in water. 

348. Mercuric Iodid, Hgl 2 . This compound is formed 
when mercury and iodin, moistened with a little alcohol, 
are rubbed together in a mortar. It is also formed when 
potassium iodid is added to a solution of mercuric chlorid. 
The precipitate thus formed is at first yellow, but soon 
becomes scarlet. 

Mercuric iodid is a scarlet-red, crystalline compound. 
It is insoluble in water, but dissolves in a solution of 
potassium iodid or mercuric chlorid, and also in alcohol. 
It melts at about 250°, and volatilizes without decompo- 
sition at a little higher temperature. 

Mercuric iodid is dimorphous, that is, it exists in two 
distinct forms: for if the scarlet crystals are heated to 
150° they become yellow; and if these yellow crystals are 
allowed to cool, and are then touched with a glass rod, 
they immediately change back to the red variety, with 
evolution of heat. 

349. Mercuric Sulfid, HgS. This important compound 
occurs in nature as the mineral cinnabar, which has a 
brownish-red color and is the principal ore of mercury. 
It is formed when mercury and sulfur, moistened with a 
little water are rubbed together in a mortar. This com- 
pound, and this method of preparing it, were known to 
the early alchemists. 



LECTURE NOTES 269 

When hydrogen sulfid is passed into a solution of a 
mercuric salt, mercuric sulfid is precipitated as a black 
amorphous powder; but if this black compound is heated, 
it volatilizes and forms the red crystals. Mercuric sulfid 
is therefore dimorphous, the red variety being crystalline, 
and the black variety amorphous. 

The red mercuric sulfid is manufactured on a large scale, 
and is known commercially as vermilion. It is used as a 
pigment and highly valued on account of its rich color. 

Mercuric sulfid is insoluble in hydrochloric, sulfuric, or 
nitric acids, but dissolves in aqua regia, forming mercuric 
chlorid. 

350. Mercuric Nitrate, *Hg(N0 3 ) 2 . This compound can 
be formed by boiling mercury, or mercuric oxid, with an 
excess of nitric acid, until sodium chlorid will no longer 
form a precipitate with a portion of the liquid (mercurous 
chlorid). 

Mercuric nitrate is a white crystalline compound. It 
dissolves in water with some difficulty unless nitric acid 
is added, and if the solution is diluted with a large excess 
of cold water, it forms a basic compound having the com- 
position 2 Hg(OH)NO s ,H 2 0, which after long boiling with 
water, forms mercuric oxid. 

This tendency to form basic compounds is a character- 
istic of mercuric salts. The other compounds of mercury 
are not important. 

THE METALS OF GROUP III. 

351. The Rare Earth Metals. The element boron, 
which occurs in the first period of Group III, is a non- 
metal, and has already been described (Page 1 73). The 



270 LECTURE NOTES 

other members of the group form two divisions, A and B, 
which are analogous to the divisions in Group II. Divi- 
sion A contains the elements scandium, yttrium, lantha- 
num, and ytterbium, which are among the rarest of all 
the elements. These four elements, together with certain 
other rare elements belonging to the preceding and suc- 
ceeding groups, are found in nature more or less closely 
associated, and so they are generally classed together and 
called the rare earth metals. The other elements are 
gadolinium, thulium, terbium, cerium, erbium, praseodym- 
ium, neodymium, and samarium. To these may also be 
added other so-called elements, such as decipium, dyspro- 
sium, and holmium, about which so little is known that 
they have not as yet been generally accepted by chemists 
as true elements. 

Most of the rare earths are so rare that very little is 
known about them and, in the case of some of them, there 
is considerable doubt as to their having been obtained free 
from other metals, and still more doubt regarding their 
position in the periodic system. Most of them have been 
inserted in the table on page 192, in the place they seem 
most likely to fill, if the atomic weights announced for 
them prove on further investigation to be correct. In 
the case of some of them it will probably be found that 
they contain other known elements, and possibly some 
elements as yet unknown. Further study of these ele- 
ments will be omitted for the present. 

352. The Aluminum Group of Metals. The remaining 
elements of Group III which belong to division B, are 
known as the aluminum group of metals, and include alu- 
minum, gallium, indium, and thallium. Of these four, 
aluminum is of considerable importance, and is the most 



LECTURE NOTES 271 

widely distributed of all known metals; thallium is a very 
rare metal; while gallium and indium are among the rarest 
of all known metals. 

These four metals bear about the same relationship to 
each other as do the metals of the magnesium group. 
They are all trivalent, and, in addition, thallium forms a 
series of compounds in which it is univalent. The metals 
are quite difficult to obtain, but when obtained are per- 
manent in the air at the ordinary temperature, but become 
oxidized when strongly heated. Thallium shows some 
curious anomalies and both as a metal and in many of its 
compounds quite closely resembles lead. 

Aluminum. Symbol, Al. Atomic Weight, 27.1. 

353. History and Occurrence. Although this metal 
was not obtained until the early part of the nineteenth 
century, some of its compounds, especially common alum, 
were known to the early alchemists, and highly valued by 
them. The metal itself was first obtained by Wohler in 
1827, by fusing together potassium and aluminum chlorid 
in a closed crucible. 

Aluminum is the third most widely distributed element 
known, and constitutes about 7.8 per cent, of the earth's 
crust. It does not occur free in nature, but its compounds 
are very numerous, occur everywhere, and in enormous 
quantities. It occurs in the mineral corundum, Al 2 O s< 
which when pure, transparent, and colored red or blue, 
forms the gems ruby and sapphire. iVn impure variety of 
corundum is called emery. It is also found in diaspore, 
AIO(OH), in bauxite, A1 2 3 , 2 H 2 0, and in cryolite, 
A1F 3 ,3 NaF. By far the largest quantity of aluminum is 
found among the silicates. It is the chief constituent of 



272 LECTURE NOTES 

the feldspars, common feldspar, or orthoclase, being 
(K,Na) 2 0,Al 2 3 ,6 Si0 2 . The decomposition of the feld- 
spars forms the clays, the purest of which is kaolin, 
Al 2 3 ,2 Si0 2 ,2 H 2 0. The micas all contain aluminum, 
common mica, or muscovite, being 2 H 2 0, K 2 0, 3 A1 2 3 , 6 Si0 2 . 
The feldspars and micas, together with quartz (Si0 2 ), 
form granite, which is an important constituent of the 
earth's crust. Finally there is hardly an important rock 
found in the earth's crust, except limestone and gypsum 
(Compare 311 and 308), which does not contain aluminum 
as an important constituent. 

354. Preparation. Until within a few years all the 
aluminum was prepared by decomposing the chlorid by 
means of sodium. This method, which was practically 
the one by which Wohler first prepared the metal, was 
improved by employing a double chlorid of aluminum and 
sodium. This was made by mixing aluminum oxid with 
sodium chlorid and carbon, and heating the mixture to 
bright redness in a stream of chlorin, the following reac- 
tion taking place: 
A1 2 3 -f 2 NaCl + 3 C + 3 CI, = 2(AlCl 8 ,NaCl) + 3 CO. 

This double chlorid was then decomposed by means of 
sodium, thus: 

AlCl 3 ,NaCl -f 3 Na = Al -f 4 NaCl. 

This process has now been entirely superseded by the 
electrolytic method. There are two or three of these 
methods, which differ only in unimportant details. By 
the Hall method, which is the principal one employed in 
the United States, aluminum oxid is dissolved in fused 
cryolite and fluorite in an iron pot lined with carbon, the 
mixture being kept melted by the current, which is a 
powerful one. A bundle of carbon rods forms the anode, 



LECTURE NOTES 273 

or positive pole, while the iron pot forms the cathode, 
or negative pole. The aluminum oxid, which is obtained 
almost entirely from the mineral bauxite, is decomposed 
by the current, the aluminum falling to the bottom of the 
pot, while the oxygen combines with the carbon of the 
anode to form carbon monoxid. The cryolite and fluorite 
remain unchanged, and so by removing the metal and 
adding more aluminum oxid, the process is made continu- 
ous. This process has greatly diminished the cost of 
production, so that aluminum is now a fairly common 
commercial product. 

355. Properties and Uses. Aluminum is a bluish- white 
metal which is capable of taking a high polish. It has a 
specific gravity of 2.58, and melts at 655°. It is one of 
the light metals, but is hard and compact, and of great 
tensile strength. If a bar of the metal is struck it gives 
a clear musical tone. It is very malleable and ductile, 
especially at temperatures between 100° and 150°, and 
can then be drawn into fine wire and beaten into very 
thin leaf. 

Aluminum is very difficult to solder but this operation 
has been successfully carried out by the use of pure zinc 
and Venetian turpentine. It can be readily welded elec- 
trically. 

Pure aluminum will not oxidize in the air at the ordi- 
nary temperature, but the impure metal soon becomes 
covered with a coating of the oxid. It is not affected by 
hydrogen sulfid. 

Its best solvent is hydrochloric acid, in which it dis- 
solves easily. It dissolves slowly in sulfuric acid, but is 
almost entirely insoluble in nitric acid, even when the 



274 LECTURE NOTES 

acid is concentrated. It dissolves also in solutions of 
sodium or potassium hydroxid. 

Aluminum is somewhat affected by salt water, and, 
especially in the presence of sodium chlorid, by such or- 
ganic acids as are often met with in cooking. It is, how- 
ever, less affected than either tin, copper, or silver, and 
so is coming to be employed quite extensively in making 
all kinds of cooking utensils; and since its salts are not 
poisonous, and the metal itself is so light, it is quite 
highly regarded for this purpose. It is not affected by 
organic secretions and so is better than silver for use in 
making surgical instruments. 

When powdered aluminum is mixed with certain me- 
tallic oxids, it can easily be set on fire and burned. In 
burning it takes the oxygen from the metallic oxid, 
reducing it to the metal, and producing great heat. This 
action may be used for reducing many oxids to the metal, 
for the welding of iron and for the production of very 
high temperatures. 

Aluminum is a good conductor of electricity, and, on 
account of its lightness, is much used for this purpose, 
and also for making physical instruments. It is used 
quite extensively for making alloys, of which the most 
important is aluminum bronze (229). 

COMPOUNDS OF ALUMINUM 

356. Aluminum Oxid, A1 2 3 . This compound occurs in 
nature as the mineral corundum, which crystallizes in 
hexagonal prisms. The gems known as ruby and sap. 
p/iire, are transparent corundum colored with chromium 
and cobalt compounds respectively. An impure corun- 
dum containing oxids of iron, is found in large quantities 



LECTURE NOTES 275 

and called emery. Corundum is very hard, and is used as 
an abrasive. 

Aluminum oxid, which is also called alumina, can be 
made by heating aluminum hydroxid, or certain other 
aluminum salts. It is a white amorphous powder, which 
after being strongly heated is insoluble in acids. 

357. The Aluminum Hydroxids. The normal hydrox- 
id, Al(OH) 3 , is found in nature as the mineral gibbsite. 
It is formed as a white gelatinous precipitate, when a 
soluble hydroxid is added to a solution of an aluminum 
salt. Another hydroxid has the composition AIO(OH), 
and is found in nature as the mineral diaspore. It can be 
prepared artificially by heating the normal hydroxid to 
300°. A third hydroxid, and the one having the greatest 
commercial importance, is found in nature as the mineral 
bauxite. It has the composition Al 2 0(OH) 4 , and generally 
contains more or less iron replacing the aluminum. It is 
the principal source of the metal. 

Aluminum hydroxid will combine with many soluble 
organic dyes, and precipitate them from a solution. 
These precipitates form a class of pigments called lakes. 
This explains the use of aluminum salts as mordants in 
dyeing. The aluminum hydroxid is precipitated on the 
cloth, and this holds the coloring matter in the fiber of 
the material. 

358. The Aluminates. Aluminum hydroxid is a very 
weak base, and does not form compounds with weak acids, 
such as carbonic or sulfurous acids, or with hydrogen 
sulfid. On the other hand aluminum hydroxid itself, in 
the presence of strong bases, acts like a weak acid, and 
forms a series of compounds called aluminates. . Thus 



276 LECTURE NOTES 

with sodium hydroxid it forms sodium aluminate, the ac- 
tion being as follows : 

Al(OH) 3 + NaOH = NaAlO, + 2 H 2 C 

A number of aluminates are found in nature. Among 
these are the minerals spinel, Mg(A10 2 ) 2 , chrysoberyl, 
G1(A10 2 ) 2 , and gahnite, Zn(A10 2 ) 2 . The soluble aluminates 
are easily decomposed by carbon dioxid or hydrogen sulfid 
forming aluminum hydroxid, thus: 

2 NaA10 2 + 3 H 2 -f C0 2 = 2 Al(OH) a + Na 2 CO s . 

Sodium aluminate is used as a mordant in calico print- 
ing and dyeing. 

359. Aluminum Chlorid, A1C1 3 . The common compound 
can be prepared by dissolving aluminum, or its hydroxid, 
in hydrochloric acid. It is a white or yellowish compound, 
which is easily soluble in water, and which crystallizes 
with six molecules of water, A1C1 8 , 6H 2 0. When these 
crystals are heated they decompose, thus: 

2 [A1C1 3 , 6 H 2 OJ= A1 2 3 + 6 HC1 + 9H 2 0. 

Anhydrous aluminum chlorid can best be prepared by 
leading perfectly dry hydrochloric acid gas over heated 
aluminum powder, and collecting the volatilized compound 
in a receiver. It is a white crystalline compound which 
fumes strongly in the air, and volatilizes without fusing 
at about 180°. The perfectly anhydrous compound is 
largely employed in the synthesis of organic compounds. 

360. Aluminum Sulfate, A1 2 (S0 4 ) 3 . This compound is 
prepared by dissolving aluminum hydroxid in sulfuric 
acid. It can also be obtained from clay, which should be 
as free as possible from iron. This is roasted in a rever- 
beratory furnace and then heated with sulfuric acid which 



LECTURE NOTES 277 

decomposes the clay forming the aluminum sulfate. It is 
also obtained in large quantities from bauxite and cryolite 
(353). 

It is a white crystalline compound which is easily solu- 
ble in water, and containing eighteen molecules of water 
of crystallization, A1 2 (S0 4 ) 3 , 18 H 2 0. It is used in many 
commercial processes. The substance largely used by the 
paper makers under the name of alum-cake, is an impure 
aluminum sulfate. 

When a solution of this salt is mixed with a solution of 
potassium sulfate, fine octahedral crystals of potassium 
aluminum sulfate are formed. These crystals have the 
composition KA1(S0 4 ) 2 , 12H" 2 0, or as it is sometimes 
written, K 2 S0 4 , A1 2 (S0 4 ) 3 ,24 H 2 0, and have been known 
tor a long time under the name of alum. 

361. The Alums. The name alum was at first applied 
to the compound just described, and this is still called 
common alum. At the present time the name is used to 
designate a series of compounds analogous to common 
alum. These compounds have the general formula 
M' M'" (S0 4 ) 2 ,12 Ii 2 0, where M' is a univalent element, 
or group, such as sodium, potassium, or ammonium, and 
M'" as a trivalent element, such as aluminum, iron (ferric), 
or chromium. The alums are all isomorphous, and crys- 
tallize in regular forms, usually octahedrons, containing 
twelve molecules of water of crystallization (Compare 328). 

In naming the alums, if the salt contains aluminum, 
only the name of the univalent element is introduced: 
thus, potassium alum is potassium aluminum sulfate, and 
ammonium alum is ammonium aluminum sulfate. If the 
alum contains potassium and no aluminum, it takes its 
name from the element replacing the aluminum: thus, 



278 LECTURE NOTES 

chromium alum, or chrome alum is potassium chromium 
sulfate. If the alum contains neither potassium nor 
aluminum, the names of both metals are used: thus, am- 
monium iron alum is ammonium ferric sulfate. 

The alums are all soluble in water. The solutions give 
an acid reaction, and have a peculiar astringent taste. 
When heated they lose all their water of crystallization, 
and if strongly heated further decompose forming oxids 
of the trivalent metal, and alkaline sulfates. The ammo- 
nium alums give off ammonia and leave the metallic oxid. 
Some of the alums have great commercial value, and are 
manufactured on a large scale. 

362. The Aluminum Silicates. These compounds are 
found in large quantities in nature usually combined with 
the alkaline silicates (353). When these compounds de- 
compose, as they do under the combined influence of the 
atmosphere and water, the alkaline silicates are dissolved 
and carried away by the water, the insoluble aluminum 
silicates remaining in the form known as clay. This clay 
when pure is white, and is then called kaolin, or porce- 
lain clay. When clay is dried and burned, it becomes 
hard and compact and is thus used for the manufacture of 
all kinds of earthen ware, from the common brick, to the 
finest porcelain. Porcelain is made hy heating a mixture 
of kaolin, feldspar, and quartz, until the feldspar melts 
and fills the pores of the clay more or less completely. 

A valuable blue pigment, formerly obtained from the 
rare mineral called lapis lazuli, is now manufactured in 
large quantities under the name of ultramarine. It is 
prepared by heating together kaolin, sodium sulfate, so- 
dium carbonate, carbon, and sulfur. This forms a green 
mass, which, by heating with sulfur, gives the desired 



LECTURE NOTES 279 

blue color. The constitution of the blue coloring matter 
is not known. 

Gallium. Symbol, Ga. Atomic Weight, 70. 

363. Gallium and its Compounds. This very rare 
element was discovered by Lecoq de Boisbaudran, in 1875, 
in some zinc-blende from Pierrefitte in the Pyrenees. It 
was discovered by means of the spectroscope, and is pecu- 
liarly interesting from the fact, that investigation showed 
it to be identical with ekaluminum, one of the three ele- 
ments predicted by Mendelejeff in 1869. 

Gallium has been prepared by electrolysis of an alkaline 
solution of its sulfate. It is a bluish-white metal which 
melts at the remarkably low temperature of 30°. It has 
a specific gravity of 5.9, and in its properties resembles 
aluminum. Its compounds have not been very carefully 
investigated on account of the great rarity of the element. 
It forms an ammonium alum. 

Indium. Symbol, In. Atomic Weight, 114. 

364. Indium and its Compounds. This very rare ele- 
ment was discovered in 1863 by Reich and Richter, by 
means of the spectroscope, in some zinc-blende from Frei- 
berg in Saxony. The metal is best obtained by fusing 
the oxid with sodium. 

Indium is a soft, silver-white metal, which has a specific 
gravity of 7.42, and melts at 176°. It is permanent in 
the air at the ordinary temperature. It resembles alumi- 
num and gallium, but differs from them in being easily 
soluble in nitric acid. Most of the common compounds 
have been formed, and it forms an ammonium alum. Its 
compounds, especially when moistened with hydrochloric 



280 LECTURE NOTES 

acid, impart an indigo-blue color to the non-luminous gas 
flame, hence the name indium. 

Thallium. Symbol, Tl. Atomic Weight, 204.1. 

365. Thallium and its Compounds. This rare element 
was discovered in 1861 by Crookes, by means of the spec- 
troscope, in the seleniferous deposit from a sulfuric acid 
manufactory in the Hartz Mts. It is found rather widely 
distributed in nature, but in very small quantities. The 
largest amount, about 17 per cent, is found in the rare 
mineral crookesite, (Cu, Tl, Ag) 2 Se. It is also found in 
many varieties of pyrite, and when these are roasted for 
the production of sulfuric acid, the thallium deposits in 
the chamber sludge, and in the chimney dust. 

The metal can be obtained from the sulfate by precipi- 
tation with zinc, or by electrolysis. It is a bluish-white 
metal, and much resembles lead. It is very soft and 
malleable. It has a specific gravity of 11.8, and melts at 
290°. It oxidizes easily in moist air, and when strongly 
heated in the air, burns with a beautiful green flame. 

Thallium forms both thallous and thallic compounds. 
Some of the thallous compounds are quite like those of the 
alkali metals, the metal itself being univalent, while both 
the hydroxid and carbonate are soluble in water and give 
an alkaline reaction. It acts as a univalent element in the 
formation of alums. In the thallic compounds the metal 
is trivalent, like aluminum, but otherwise shows very 
little resemblance to that metal. Thallium compounds 
color the non-luminous gas flame a bright green. All the 
soluble compounds of thallium are poisonous. 

THE METALS OF GROUP IV. 

366. The Divisions of the Fourth Group. The ele- 
ments which constitute the fourth group are somewhat 



LECTURE NOTES 281 

peculiar, the two divisions being much more closely rela- 
ted than those of the preceding groups. The typical 
element of the group is carbon, which is a non-metal, and 
does not exhibit any metallic properties in any of its 
compounds. Silicon is also a non-metal and very closely 
resembles carbon. In its elementary state silicon resem- 
bles the metals, but in its compounds it is always non- 
metallic. 

The other elements of the group are more basic, their 
metallic properties increasing with their atomic weights. 
There are no marked differences in the two divisions. The 
metals generally regarded as belonging to division A are 
titanium, zirconium, cerium, and thorium, while germa- 
nium, tin, and lead constitute division B. Excepting tin 
and lead, they are all very rare elements. The elements 
of this group are all normally quadrivalent, although 
they nearly all form compounds in which they are bivalent. 

Titanium. Symbol, Ti. Atomic Weight, 48.1. 

367. Occurrence, Preparation, and Properties. In 1789 
Gregor discovered a new element in the mineral ilmenite, 
and in 1795 Klaproth discovered a new element in the 
mineral rutile, to which he gave the name titanium. Sub- 
sequent investigation proved that the two elements were 
identical. Titanium does not occur free in nature. In 
the form of the oxid, Ti0 2 , it is found in the minerals 
rutile, octahedrite, and brookite, three minerals which have 
the same composition, but differ in crystalline form. As 
an iron titanate, it is found in the mineral i7raemte,FeTi0 3 , 
and as a calcium titano-silicate, in the mineral titanite, or 
sphe?ie, CaTiSi0 5 . It also occurs in small quantities in 
the form of the oxid, replacing silica in many silicates. 

The metal is very difficult to obtain pure, and it is 



282 LECTURE NOTES 

somewhat doubtful if it has ever been obtained pure, ow- 
ing to the ease with which it forms compounds with car- 
bon and nitrogen. It is usually prepared by reducing the 
oxid by means of sodium or magnesium, or by means of 
carbon in an electric furnace. 

It is a very hard, white metal, having a specific gravi- 
ty of 4.87. It dissolves in the common acids, especially 
sulfuric acid, and decomposes water at a high temperature. 

368. The Compounds of Titanium. In its compounds 
titanium much resembles silicon, and it is more non-me- 
tallic than metallic. Its most important compound is tit- 
anium dioxid, TiO a , which occurs in nature as given above. 

Titanium forms a series of acids analogous to the silicic 
acids (214). It also forms compounds with each of the 
halogen elements. 

Titanium is remarkable for the ease with which it com- 
bines directly with nitrogen and carbon. Normal tita- 
nium nitrid, Ti 3 N 4 , is a copper-colored compound which 
was at first supposed to be metallic titanium. Another 
nitrid, TiN, is a yellow substance, having a specific grav- 
ty of 5.18, and is as hard as the diamond. When 
titanium dioxid is fused with carbon in an electric furnace, 
titanium carbid, TiC, is formed. This carbid does not de- 
compose water at any temperature (Compare 310). It is 
often found in the cast iron prepared from ores containing 
titanium. 

Zirconium. Symbol, Zr. Atomic Weight, 90.4. 

369. Zirconium and its Compounds. This element was 
discovered by Klaproth, in 1789, in the mineral zircon. 

Zirconium does not occur free in nature, and its most 



LECTURE NOTES 283 

important compound is the mineral zircon, ZrSi0 4 , which 
occurs very widely distributed, as a very small but im- 
portant constituent of most of the eruptive rocks, which 
form the greater portion of the earth's crust. It is also 
found in a few other rare minerals. 

Zirconium, like carbon, exists in three allotropic modi- 
fications. When potassium zirconofluorid, K 2 ZrF 6 , is 
heated with potassium, the metal is obtained as an amor- 
phous black powder, which readily takes fire on being 
warmed in the air. When zirconium oxid, Zr0 2 , is heated 
in an electric furnace with carbon, the metal is obtained 
as a very hard, steel-gray, crystalline substance, which 
has a specific gravity of 4. 25. A form analogous to graph- 
ite has also been isolated. 

The compounds of zirconium are mostly unimportant, 
and resemble those of titanium. Zirconium oxid, Zr0 2 , has 
recently become of value as an illuminating material. 
For this purpose it is pressed into tablets, which, when 
heated in the oxyhydrogen flame, emit an intensely white 
light, that is superior in quality to the Drummond lime 
light (302), and is much employed in optical research 
work. 

Cerium. Symbol, Ce. Atomic Weight, 139. 

370. Cerium and its Compounds. This element is 
always found associated with the rare earths of the third 
and fifth groups, although being quadrivalent, it undoubt- 
edly belongs in this fourth group. It was discovered in 
1803 in the mineral cerite, a silicate of some of the rare 
earths, the exact formula of which is doubtful. 

The metal was first prepared by heating the chlorid 
with sodium, and may be obtained by the methods em- 



284 LECTURE NOTES 

ployed for the other metals of this group. It is a soft, 
gray metal which is quite malleable. It has a specific 
gravity of 6.7. It takes fire more easily than magnesium, 
and burns with a very brilliant light. 

Cerium forms both cerous, and eerie compounds. In the 
cerous compounds the metal is trivalent, and in the eerie 
compounds it is quadrivalent. 

Thorium. Symbol, Th. Atomic Weight, 232.6. 

371. Thorium and its Compounds. This element was 
discovered in the mineral thorite in 1828, by Berzelius. 
The mineral thorite, a variety of which is called ora?igeite, 
is essentially thorium silicate, ThSi0 4 . It is also found 
in a number of rare minerals especially in monazite, 
(Ce, La)P0 4 , which usually contains thorium oxid, the 
amount varying from mere traces, to as high as 18 per 
cent. The monazite found in North Carolina, and in 
Brazil, constitutes the principal source of thorium com- 
pounds at the present time. 

The metal is obtained by heating the double chlorid of 
potassium and thorium with sodium in an iron cylinder. 
It is a white, crystalline metal, having a specific gravity 
of 11. It takes fire when heated in the air, and burns 
with a bright flame. 

The salts of thorium resemble those of the other mem- 
bers of this group, and, excepting the oxid, have no 
commercial value. 

372. The Welsbach Light. Thorium oxid, or thoria, 
Th0 2 , is even more valuable for illuminating purposes 
than zirconium oxid, since it becomes highly incandescent 
at the temperature of the Bunsen flame. This property 



LECTURE NOTES 285 

has been utilized by Auer von Welsbach for the construc- 
tion of the incandescent gas lamp which bears his name. 
A finely meshed mantle of cotton, or linen, is saturated 
with a solution of thorium nitrate, containing a small 
amount of cerium nitrate. This is made to surround a 
non-luminous gas flame, or Bunsen burner. On being 
ignited, the vegetable material is burned away, leaving a 
mantle of thorium and cerium oxids, which retains the 
form of the original mantle, and emits a bright, soft light. 
The mantles contain between 98 and 99 per cent, of tho- 
rium oxid. A new mantle gives a light of from 75 to 100 
candle power, and, unless broken by accident, will last 
about 1000 hours. The candle power of the light dimin- 
ishes quite rapidly at first, and then more gradually, 
until at the end of the life of an ordinary mantle, its ini- 
tial candle power has often been reduced one half. This 
is probably owing to the slow volatilization of the oxids. 

Germanium. Symbol, Ge. Atomic Weight, 72.5. 

373. Germanium and its Compounds. This element 
was discovered in 1886, by Winkler, while investigating 
a silver mineral recently discovered at Freiberg in Saxo- 
ny. This mineral, which is called argyrodite, is a sulfo- 
germanate of silver, and has the composition 4 Ag 2 S,GeS 2 . 
It is also found in minute quantities in the very rare 
minerals canfieldite, and euxenite. 

Germanium is one of the rarest of all known elements, 
but is interesting from the fact, that when discovered, it 
proved to be the ekasilicon predicted by Mendelejeff many 
years before. It forms two classes of salts, germanous 
and germanic. In the former it is bivalent, and in the 
latter quadrivalent. None of its salts have any com- 
mercial value. 



286 LECTURE NOTES 

Tin. Symbol, Sn. (Stannum). Atomic Weight, 119. 

374. History and Occurrence. Tin is one of the metals 
known to the ancients, and is mentioned in the Pentateuch. 
It is often mentioned by Pliny in his Natural History under 
the name pi umbum album, to distinguish it from plum- 
bum nigrum, or lead. The Greek alchemists at first called 
it Hermes, but later it received the name of Jupiter, and 
was referred to by the sign of the planet Jupiter, Q)., in 
their writing. 

Tin has been found in the native state in minute quan- 
ties in a few localities. Its most important ore is the 
mineral cassiterite, or tinstone, Sn0 2 , this ore being practi- 
cally the only source of tin. The few other minerals 
which contain it have no commercial value. 

The oldest and best known tin mines in the world are 
those of Cornwall in England, where tin has been ob- 
tained from the time of the Phoenicians up to the present. 
Other important localities where tin is obtained are the 
Malay Peninsula, the Island of Bancain the Malay Archi- 
pelago, Australia, Bolivia, and a few European localities. 
There are a few localities in the United States where tin 
has been found, but up to the present time the mines have 
not proved profitable, and most of them have been aban- 
doned. About one-half of the entire quantity comes from 
the native state of Perak in the Malay Peninsula. 

375. Preparation. The process of reducing tin ores 
is a simple one. The ore, after being separated as much 
as possible from the gangue, is roasted in a furnace of 
peculiar construction, in order to remove sulfur and arsen- 
ic, which are very commonly found in small quantities in 
connection with tin ores. The furnace consists of a long 
iron cylinder, lined with fire brick, and placed in an in- 



LECTURE NOTES 287 

clined position. The fire is at the lower end, while the 
upper eod is connected with condensing chambers. The 
ore is brought into the upper end of the cylinder, which 
is made to revolve, so that the ore gradually works down 
the cylinder, and falls into a place prepared for it, while 
the arsenic and sulfur volatilize and pass into the con- 
densing chambers. 

The roasted ore is then mixed with coal and smelted 
in a reverberatory furnace, after which the metal is 
drawn off and cast into ingots. The ingots are placed in 
another furnace and carefully heated, when the easily 
fusible tin runs off, leaving the less fusible impurities, 
which consist of iron and copper oxids, and alloys of tin 
and arsenic. The tin is further purified by being melted 
and stirred with pieces of green wood. This causes a scum 
to rise, which carries with it the most of the remaining 
impurities. 

376. Properties. Tin is a bright, white metal, which 
does not oxidize in either moist or dry air, It has a spe- 
cific gravity of 7.3, and melts at 232°. It is harder than 
lead but can be cut with a knife. It has a crystalline 
structure, and when bent, it emits a faint crackling sound, 
known as the "cry of tin.'' It is quite malleable, and can 
be rolled into thin sheets known as tin-foil, but at 200°, 
or at very low temperatures, it becomes brittle, and can 
then be powdered. 

Tin dissolves easily in hydrochloric acid, with evolution 
of hydrogen and the formation of stannous chlorid. Con- 
centrated sulfuric acid also dissolves it with evolution of 
sulfur dioxid. Cold dilute nitric acid dissolves it with 
the formation of stannous nitrate, but the concentrated 
acid converts it into metastannic acid. 



288 LECTURE NOTES 

Tin forms two classes of compounds called stannous and 
stannic. In the former it is bivalent and in the latter, 
quadrivalent. The stannous compounds are all easily- 
oxidized to stannic compounds. 

377. Uses. Tin is employed for a large number of 
purposes, such as the making of househeld untensils, the 
manufacture of tin-foil, the tinning of easily oxidized 
metals, and the preparation of alloys. Tinning consists 
in coating other metals, with a thin film of tin, by dip- 
ping them into a bath of the molten metal. Tin plate, 
which is so extensively employed for roofing and for 
making household utensils, is tinned sheet-iron. 

The alloys of tin are of grea\ importance. Tin and 
lead will mix in all proportions. Pewter is made of three 
or four parts of tin and one of lead. Common solder is 
made of equal parts of tin and lead, while fine solder con- 
tains two parts of tin and one of lead, and coarse solder, 
one part of tin and two of lead. Most of the other impor- 
tant alloys of tin have already been mentioned (229). 

Tin readily unites with mercury to form an amalgam 
which contains about 50 per cent, of tin. It is used for 
coating the glass used for mirrors. 

THE STANNOUS COMPOUNDS. 

378. Stannous Chlorid SnCl 2 . This compound is pre- 
pared on a large scale, by dissolving tin in hydrochloric 
acid. When this solution is evaporated, crystals having 
the composition, SnCl 2 , 2 H 2 are formed. When tin is 
heated in dry hydrochloric acid gas, the anhydrous stan- 
nous chlorid is formed. This melts at 250° and distils 
without decomposition at 606°. 

Stannous chlorid is easily soluble in a small quantity 



LECTURE NOTES 289 

of water, but with an excess of water, and especially 
when exposed to the air, a basic chlorid separates, thus: 

3 SnCl 2 -f O + H 2 == 2 Sn(OH)Cl + SnCl 4 . 

Stannous chlorid is a powerful reducing agent, and 
readily acts upon compounds containing either oxygen or 
chlorin; thus it reduces mercuric chlorid, first to mercu- 
rous chlorid, and then to metallic mercury. Under the 
name of tin salt it is used as a mordant in dyeing. 

379. Other Stannous Compounds. Stannous hydroxid, 
Sn(OH) 2 , is formed when a solution of a stannous salt is 
acted upon by a soluble hydroxid or carbonate. It is a 
white compound, which, if carefully heated in an atmos- 
phere of carbon dioxid, forms black stannous oxid, SnO. 

Stannous sulfide SnS, is a dark brown amorphous com- 
pound, which may be formed by leading hydrogen sulfid 
into a stannous solution, or by fusing together tin and 
sulfur. 

There are other stannous compounds known, but they 
are of little importance. 

THE STANNIC COMPOUNDS. 

380. Stannic Oxid, Sn0 2 . This compound is found in 
nature as the mineral cassiterite or tin-stone, occurring 
both in crystals and in compact masses. These are usu- 
ally colored dark brown, or black, and have a specific 
gravity of 6.8. It can be prepared artificially by strongly 
heating tin in the air, or by adding ammonia to a stannic 
solution and igniting the precipitate thus formed. 

Stannic oxid, when prepared in this way, is a yellowish- 
white amorphous powder, which is known commercially 



290 LECTURE NOTES 

as putty powder, and is much used for polishing stone and 
glass. 

381. The Stannic Acids. The hydroxids correspond- 
ing to stannic oxid, are known as the stannic acids, of 
which there are two. These might be expected to corre- 
spond to the two principal silicic acids, and so have the 
composition H 2 Sn0 3 and H 4 Sn0 4 . As a matter of fact 
both acids exist in all degrees of hydration, between the 
limits of these two formulas, and their differences are 
seen in their physical properties, rather than in their 
chemical composition. They are known as stannic and 
metastannic acids, and also as oc-stannic and fi -stannic 
acids. 

Stannic acid, or a-stannic acid, H 2 SnO s , is formed when 
calcium carbonate acts upon a solution of stannic chlorid, 
thus: 
2 CaC0 3 + SnCl 4 + H 2 = 2 CaCl 2 + 2 C0 2 + H 2 SnG 3 . 

It is a white compound which is slightly soluble in wa- 
ter and, when dried in a vacuum has the composition 
H 2 SnO s , but when dried in the air, has the composition 
H 2 SnO s , H 2 0, or H 4 Sn0 4 . It forms a number of salts, the 
sodium salt being the most important one. This has the 
composition Na 2 Sn0 3 , 3 H 2 0, is soluble in water, and used 
as a mordant in dyeing. 

Metastannic acid, or p-stannic acid, has commonly the 
composition H 2 Sn 5 O n , 4 H 2 0, or H 10 Sn 5 O 15 . It is formed 
as a white insoluble powder when tin is acted upon by 
concentrated nitric acid. Its salts are quite complicated, 
its sodium salt having the composition Na 2 Sn 6 O n , 4 H 2 0, 
which is also written Na 2 Sn0 3 , 4 Sn0 2 , 4 H 2 0. 

382. Other Stannic Compounds. Stannic chlorid,SnCl^ 
is prepared by passing a stream of dry chlorin over melted 



LECTURE NOTES 291 

tin in a glass retort, or by the oxidation of stannous chlo- 
rid. It is a colorless, fuming liquid which boils at 114°, 
and was known to the alchemists as the fuming liquid of 
Libavius. It dissolves easily in water, but if a dilute 
solution is boiled, it decomposes, forming metastannic 
acid. It is used as a mordant in dyeing. 

Stannic sulfd, SnS 2 , is formed by leading hydrogen 
sulfid into a stannic solution. It is a bright yellow com- 
pound which is used as a pigment, and called mosaic gold. 

Tin forms sulfo-stannates by the action of the alkaline 
sulfids upon stannous or stannic sulfids. They are anal- 
ogous to the corresponding compounds of arsenic (Compare 
166). The other stannic compounds are not important. 

Lead. Symbol, Pb. (Plumbum). Atomic Weight, 206.9. 

383. History and Occurrence. Lead is one of the metals 
which was known to the ancients. It is mentioned in the 
Bible, as well as in the works of Homer, and of Pliny, the 
latter being the first to point out a distinction between 
lead and tin (Compare 374). The alchemists assigned it 
to the planet Saturn, with which it was supposed to have 
some connection, and it is always referred to in their 
writings by the symbol of Saturn, viz. : j^. 

Lead has been found, in small quantities, free in nature, 
in certain volcanic regions. Its most important ore is 
the mineral galenite or galena, PbS, which occurs very 
widely distributed, and often in enormous quantities. 
Lead occurs in a very large number of minerals, among 
the more important of which may be mentioned, cerussite, 
PbC0 3 , anglesite, PbS0 4 , matlockite, PbCl 2 , PbO, and 
pyromorphite, 3 Pb 3 (P0 4 ) 2 , PbCl 2 . Nearly all the lead in 
the world is obtained from galena. 



292 LECTURE NOTES 

384. Preparation. The process of preparing lead from 
galena is a simple one. The ore is first partially roasted 
in a reverberatory furnace, whereby some of the lead sul- 
fid is oxidized to lead sulfate, and lead oxid, thus : 

PbS + 2 2 = PbS0 4 . 
2 PbS + 3 O, = 2 PbO + 2 S0 2 . 

When a sufficient quantity of the galena has thus been 
oxidized, the temperature is raised, whereupon the lead 
sulfate and oxid react upon the unchanged sulfid, forming 
metallic lead and sulfur dioxid, thus: 

PbS0 4 -f PbS = 2 Pb + 2 S0 2 
2 PbO + PbS = 3 Pb + S0 2 

Lead may also be obtained from galena by fusing with 
iron, thus: 

PbS + Fe = FeS 4- Pb. 

Lead nearly always contains some silver, and generally 
other impurities. The silver is removed by one of the 
processes for desilverizing lead which have already been 
given under silver (See 288). The other impurities are 
removed in various ways, the method depending upon 
their nature, and the purpose for which the lead is to be 
used. 

385. Properties and Uses. Lead is a bluish-white metal 
which is crystalline in structure. It is so soft that it is 
easily cut with a knife, and will leave a streak when rubbed 
upon paper. The freshly cut surface has a bright luster, 
but on exposure to the air, it soon becomes dull gray 
from surface oxidation. It has a specific gravity of 11.37 
and melts at 326°. When it is strongly heated in the 
air, it burns, forming lead oxid, PbO. 



LECTURE NOTES 293 

Lead is insoluble in dilute hydrochloric and sulfuric 
acids, and only very slightly soluble in the concentrated 
acids. It is easily soluble in nitric and acetic acids. Lead 
is not affected by pure water in the absence of air, but in 
contact with air, lead hydroxid, Pb(OH) 2 , is formed, which 
is slightly soluble in water. Water containing small 
quantities of phosphates and carbonates has no solvent 
action upon lead. On the other hand, water containing 
nitrates (rain water) dissolves lead to quite an extent. 
All the soluble compounds of lead are poisonous. 

Lead is an extremely useful metal and employed in 
numberless ways, both on account of the ease with which 
it can be worked, and its power of resisting the action of 
so many corroding substances. It is used in the form of 
sheet lead and as lead pipe, and is a constituent of many 
useful alloys (Compare 229 and 377). 

COMPOUNDS OF LEAD. 

386. The Oxids of Lead. Lead forms five compounds 
with oxygen. These are: 

Lead suboxid, Pb 2 0, 

Lead monoxid, or Lead oxid, PbO, 

Lead sesquioxid, Pb 2 3 , 

Red lead, or Minium, Pb 3 4 , and 

Lead dioxid, or Lead peroxid, PbO a . 

The most important one is lead oxid, which is a strong- 
ly basic compound, in which the metal is bivalent. Lead 
dioxid is a weak base, and, in the presence of the alkaline 
bases, acts as a weak acid. Lead suboxid is a black pow- 
der, and is of little importance. The other two oxids, of 
which red lead only is of importance, may be considered 



294 LECTURE NOTES 

as salts formed by the combination of the basic lead oxid, 
PbO, with the acid lead dioxid, Pb0 2 . 

387. Lead Monoxid, or Lead Oxid, PbO. This com- 
pound was known to the ancients. It is formed when 
lead is strongly heated in the air, and so is obtained in 
large quantities by the cupellation of lead, in the process 
of extracting silver (288). It is also formed when certain 
lead salts are heated, and is the final product obtained by 
heating any of the other oxids of lead. 

Lead oxid is a yellow amorphous powder, which is 
known commercially as massicot. It melts at a high tem- 
perature, and, on cooling, forms a crystalline mass known 
as litharge. It is used in the manufacture of flint glass, 
and as a glaze for earthen ware. 

A corresponding lead hydroxid, Pb(OH) 2 , is formed 
when a soluble hydroxid is added to a solution of a lead 
salt. 

388. Red Lead, Pb s 4 . This compound was also known 
to the ancients under the name minium. It is prepared 
by heating lead oxid for a long time in the air, at a tem- 
perature of about 400°. It is a bright red powder, which 
has a specific gravity of 8. 6. It is generally regarded as 
a compound of lead oxid with lead dioxid, 2 PbO, Pb0 2 , 
or as a basic lead plumbate, having the composition 

o<PbZS> Pb =°- 

When strongly heated, it gives up a portion of its oxy- 
gen and forms lead oxid. It is largely used as a pigment. 

389. Lead Dioxid, or Lead Peroxid, Pb0 2 . This com- 



LECTURE NOTES 295 

pound may be prepared in many ways. A common method 
is to oxidize red lead with dilute Ditric acid, thus: 

Pb 3 4 -J- 4 HN0 3 = Pb0 2 + 2 Pb(N0 3 ) 2 + 2 H 2 0. 

The other methods of preparation are analogous, and 
consist in oxidizing lead oxid by means of some strong 
oxidizing agent. 

Lead dioxid is a heavy dark-brown powder, and is often 
called the puce-colored oxid of lead. It loses a portion of 
its oxygen when exposed to sunlight, forming red-lead, 
and on strongly heating forms lead oxid. It is therefore 
an oxidizing agent. It acts as an acid-forming oxid in 
the presence of strong bases, and forms thus a series of 
salts called plumbates. The corresponding acids are 
known as plumbic and metaplumbic acids, and are analo- 
gous to the stannic acids. 

390. Lead Chlorid, PbCl 2 . This compound is obtained 
when hydrochloric acid, or a soluble chlorid, is added to 
a solution of a lead salt ; or when lead oxid or carbonate 
is dissolved in hydrochloric acid. 

It is a white crystalline compound which is only slight- 
ly soluble in water. When it is heated in the air it forms 
an oxy-chlorid, PbCl 2 , PbO, which is identical in compo- 
sition with the mineral matlockite. It forms a number of 
double salts with other metallic chlorids. 

A lead tetrachlorid, PbCl 4 , is also known. It is a 
yellow, fuming liquid, which is easily decomposed by con- 
tact with water, forming lead chlorid, and chlorin. 

391. Lead Sulfid, PbS. This compound is found in 
nature in large quantities, and widely distributed, as the 
mineral galena. In this form it is a bluish-gray substance 



296 LECTURE NOTES 

which crystallizes in cubes. It has a bright metallic lus- 
ter, and a specific gravity of about 7.5. 

When lead is heated in sulfur vapor, lead sulfid is 
formed as a gray crystalline mass, which closely resembles 
the native galena. When hydrogen sulfid is passed into 
a solution of lead nitrate, lead sulfid is formed as a black 
amorphous powder. 

When lead sulfid is heated in the air, it oxidizes to lead 
sulfate. Dilute nitric acid dissolves it, forming lead 
nitrate, but the concentrated acid oxidizes it to lead 
sulfate. 

392. Lead Sulfate, PbS0 4 . This compound occurs in 
nature in transparent crystals as the mineral anglesite. 
It is formed as given in the preceding paragraph, and by 
the action of hot concentrated sulfuric acid on lead. It 
is also formed when sulfuric acid, or a soluble sulfate, is 
added to a solution of a lead salt. 

Lead sulfate is a white crystalline compound. It is 
only very slightly soluble in water, but dissolves to a 
considerable extent in concentrated sulfuric acid. It also 
dissolves in a solution of sodium hydroxid, or ammonium 
acetate. This distinguishes it from barium sulfate, which 
it otherwise resembles (317). When ignited with carbon 
it is reduced to the sulfid. 

393. Lead Nitrate, Pb(N0 3 ) 2 . This compound was known 
to the early alchemists. It is formed when lead, or its 
oxid or carbonate is dissolved in nitric acid. When the 
solution is evaporated, the salt crystallizes out in regular 
octahedrons, which contain no water of crystallization. 

Lead nitrate is a white compound, which is quite easily 
soluble in water. It is the only inorganic salt of lead 



LECTURE NOTES 297 

that is soluble in water to any extent, and so finds con- 
siderable use in the laboratory. It is used in the manu- 
facture of lead chromate, which is used as a pigment 
under the name of chrome yellow. It is also used in calico 
printing and dyeing. 

394. Lead Carbonate, PbC0 3 . This compound is found 
in nature as the mineral cerussite. It is formed when a 
cold solution of sodium or ammonium carbonate, is added 
to a solution of a lead salt. 

The normal lead car-bonate has no commercial value, 
but lead forms several basic carbonates, one of which is 
of great commercial importance. This compound has the 
composition 2 PbC0 3 , Pb(OH) 2 , and is known commercially 
as white lead. There are several processes by which this 
substance is manufactured on a large scale, and which 
differ only in technical details. 

The oldest process, and one which is still employed in 
many parts of the world, is known as the Dutch process. 
By this process, lead plates are rolled up into spirals, and 
placed on perforated wooden shelves, in earthen ware pots 
of peculiar construction. Common vinegar, or dilute 
acetic acid, is poured into the lower part of these pots, 
which are then placed side by side on a thick layer of 
spent tan-bark, on the floor of a shed, and covered with 
boards. Instead of the tan-bark, horse-dung was at first 
used, and is still in use in certain localities. Another 
layer of the tan-bark is placed on the boards, and on this 
another series of pots, and so on until the shed is filled. 
The whole is then covered with a thick layer of the bark, 
and allowed to stand for several weeks. 

The heat caused by the fermentation of the tan-bark, or 
dung, causes the acid to evaporate, and this dissolves the 



298 LECTURE NOTES 

lead, forming basic lead acetate, while at the same time 
the carbon dioxid, generated by the fermentation, converts 
this basic lead acetate into basic lead carbonate, and liber- 
ates the acetic acid. The liberated acid at once dissolves 
more lead, and the process is repeated, so that a small 
amount of the acid is able to convert a large amount of 
the metal into white lead. 

By the French method, which is in use not only in France, 
but also in many other countries, the basic lead acetate is 
made by dissolving lead oxid in acetic acid, and the basic 
carbonate precipitated from the solution by means of car- 
bon dioxid. 

White lead is a heavy amorphous powder. It is very 
poisonous, and becomes blackened when exposed to hydro- 
gen sulfid. In spite of these disadvantages, white lead is 
almost universally employed as a pigment in the manu- 
facture of paint. It is not infrequently adulterated with 
permanent white, BaS0 4 (317), a small quantity of which 
does not seem to do any harm. For some purposes zinc 
white, ZnO, is a more desirable pigment, as it is not 
blackened by hydrogen sulfid (333). 

THE METALS OF GROUP V. 

395. General Characteristics. The elements which 
belong to this group, like those which belong to Group 
IV, resemble one another very closely in their chemical 
properties, although they differ in some other respects. 
Three of the elements, nitrogen, phosphorus, and arsenic, 
are well defined non-metals, and have already been de- 
scribed. The other elements, while much more basic than 
these three, still have many non-metallic properties. 
Those of the metals which are generally regarded as be- 



LECTURE NOTES 299 

longing to division A, are vanadium, columbium, praseo- 
dymium, and tantalum, while antimony, erbium and 
bismuth, constitute division B. Nitrogen is the typical 
element of the group, while phosphorus and arsenic are 
perhaps a little more closely related to the elements of 
division B (Compare 114). 

Praseodymium and erbium are usually found with the 
so-called rare earths (351); and while they are usually 
assigned to Group V, they are such rare elements, that 
too little is known about them to make it certain that 
they really belong to it. 

Each of the members of this group forms a series of 
compounds with oxygen, the valence of which varies from 
two to five. All of the higher oxids, and many of the 
lower ones, are acid-forming. Several of the elements 
yield a compound with hydrogen in which they are 
trivalent. 

Vanadium. Symbol, V. Atomic Weight, 51.5. 

396. History and Occurrence. The existence of this 
rare element was first made known in 1801, by Del Rio, 
who obtained some of its compounds from a lead ore found 
in Mexico, and who named the new element erythronium. 
So little was known of it that it was generally believed to 
be an oxid of chromium until 1830, when Sefstrom again 
discovered it in some Swedish iron, and called it vana- 
dium. In the same year Wohler proved that the two 
elements were the same, and in 1867 Roscoe succeeded in 
obtaining the metal. 

Vanadium is found in only a few rare minerals, the 
most important of which are vanadinite, 3 Pb 3 (V0 4 ) 2 , 
PbCl 2 , and mottramite, (Pb,Cu) 3 (V0 4 ) 2 ,2 (Pb,Cu)(OH) 2 . 



300 LECTURE NOTES 

It is also found in certain iron ores, and, in traces, in a 
number of rocks and soils. 

397. Preparation, Properties and Compounds. The 

metal is prepared by beating vanadium dichlorid in a 
stream of perfectly pure, dry hydrogen. 

It is a grayish- white, crystalline metal, with a bright 
metallic lustre. It is fairly permanent in the air at the 
ordinary temperature, and does not decompose water; 
but when heated in the air it takes fire easily and burns 
brilliantly. 

Vanadium forms five compounds with oxygen, all the 
higher ones being acid-forming oxids. Its acids are anal- 
ogous to those of phosphorus and arsenic. Orthovanadic 
acid, H 3 V0 4 , does not exist free, but its salts are quite nu- 
merous. In its basic relations vanadium forms compounds 
with the halogens, and with most of the common acids. 
It forms a compound with nitrogen by direct union of the 
elements. 

Columbium. Symbol, Cb. Atomic Weight, 93.7. 

398. History, Occurrence and Properties. This element 
is quite closely related to tantalum, the two being nearly 
always found together. It was discovered in 1801, by 
Hatchett, in the mineral columbite from Haddam, Conn., 
and named by him columbium. For a long time after the 
discovery of tantalum, the following year, the two ele- 
ments were believed to be identical. In 1844, Rose ob- 
tained the oxid of a supposed new element which he called 
niobium, and which was afterwards found to be identical 
with columbium. During the few years following, a 
number of supposed new elements were discovered. 
Among these the elements called pelophim and diauium 



LECTURE NOTES 301 

were found to be identical with columbium, and ilmenium 
and neptunium were found to be mixtures of columbium 
and tantalum. 

Columbium is found in several rare minerals the more 
important of which are columbite, (Pe,Mn)(Cb0 3 ) 2 , and 
pyrochlore, (Ca,Fe)(Cb0 3 ) 2 , Ce(Ti, Th)0 3 , NaF, the others 
being principally columbates and tantalates of the rare 
earths. 

The metal may be prepared by reducing the chlorid 
with hydrogen. It is a steel-gray metal, having a speci- 
fic gravity of 7, and takes fire at a low temperature when 
heated in the air. It dissolves in concentrated sulfuric 
acid. In its compounds, which resemble those of vana- 
dium, the element acts like a non-metal, being usually 
found in the acid radical. 

Columbium is frequently called niobium, but, as its 
original discoverer called it columbium, this name is gen- 
erally preferred, especially by American chemists. 

Praseodymium. Symbol, Pr. Atomic Weight, 140.5. 

399. History and Compounds. In 1839, Mosander, 
while investigating cerium oxid, succeeded in separating 
from it another oxid containing a new metal, which he 
called lanthanum. In 1844, he discovered that this new 
element was not pure, but contained another element, 
which he called didymium. In 1879, Lecoq de Boisbau- 
dran separated from didymium another new element, 
which he called samarium. In 1885 Welsbach again re- 
solved didymium into two elements, which he called 
praseodymium and neodymium. It is now believed that 
the latter still contains at least one as yet unknown ele- 



302 LECTURE NOTES 

merit. Not more than one or two of these elements have 
been isolated, and they in a doubtful state of purity. 

The salts of praseodymium are green, and form a green 
solution, while those of neodymium are amethyst color, 
or pink. 

Tantalum. Symbol, Ta. Atomic Weight, 182.8. 

400. Tantalum aDd its Compounds. Some of the his- 
tory of this element was given under columbium (398). 
The mineral called tantalite, has essentially the composition 
(Fe, Mn)(Ta0 3 ) 2 , and is very closely analogous to colum- 
bite. In nearly all the minerals containing columbium, 
tantalum is found to be present in varying quantities. 

The element was discovered in 1802, by Ekeberg, in a 
mineral which he called yttrotantalite, and which is a tan- 
talate and columbate of iron, calcium, and some of the 
rare earths. 

When potassium tantalofluorid, K^TaF 7 , is ignited with 
potassium, a black powder is obtained which is largely 
metallic tantalum, although it is doubtful if this element 
has ever been obtained very pure. 

Tantalum does not dissolve in any single acid, except 
hydrofluoric acid, nor does it dissolve in aqua regia. Its 
compounds are very similar to those of columbium, the 
element nearly always acting as a non-metal. 

Antimony. Symbol, Sb (Stibium). Atomic Weight, 120.4. 

401. History and Occurrence. The mineral Stib- 
nite, Sb 2 S 3 , which is the principal ore of antimony, has 
been known since very early times, having been employed 
by women in the East for painting the eyebrows. It is 
mentioned in the Bible, and was known to the early Greek 



LECTURE NOTES 303 

philosophers. The preparation of the metal, which was 
known to the alchemists by the names stibium and anti- 
?nonium, was first described by Basil Valentine in the 
15th century. A number of compounds of antimony were 
undoubtedly made about this time, and these, as well as 
the metal itself, were highly valued by the alchemists, 
and especially by the later iatro-chemists. 

Antimony is a fairly common element. It is sometimes 
found free in nature, but much more commonly in combi- 
nation with sulfur, or with sulfids of copper and nickel. 
There are many minerals which contain antimony, but by 
far the most important of these is stibnite, Sb 2 S 3 , from 
which the metal is nearly always obtained. This is 
found in nearly all parts of the world. The other anti- 
mony minerals have little or no commercial importance. 

402. Preparation. The process by which antimony is 
obtained from stibnite is a simple one. In order to sepa- 
rate the mineral from the surrounding rock impurities, it 
is generally fused in iron cylinders having openings at the 
bottom through which the fused material may run out. 
The fused sulfid, or the original ore, if it was sufficiently 
pure, may be fused with metallic iron, when the ore is 
reduced, thus: 

Sb 2 S 3 + 3 Fe = 3 FeS + 2 Sb. 

The fused sulfid may also be roasted in a reverberatory 
furnace, and thus changed to the oxid, Sb 2 4 . The oxid is 
reduced with carbon, or with crude tartar, in earthen 
crucibles, thus: 

Sb 2 4 + 4 C = 2 Sb + 4 CO. 

If the metal is required pure, it is afterwards refined. 



304 LECTURE NOTES 

403. Properties and Uses. Antimony is a brilliant, 
bluish-white metal, having a leafy, crystalline structure. 
It is very hard, and so brittle that it is easily powdered 
in a mortar. It has a specific gravity of 6.7 and melts at 
630°. It is not altered in the air at the ordinary temper- 
ature, but when heated it burns to the oxid. It volatil- 
izes in the air at a red heat, and may be distilled in a 
current of hydrogen at about 1600°. 

Pure antimony is not acted upon by hydrochloric acid, 
or by dilute sulfuric acid, although the common metal, 
which usually contains more or less impurity in the form 
of copper, lead or arsenic, dissolves quite easily in con- 
centrated hydrochloric acid. Nitric acid converts antimony 
into a white powder, the composition of which varies with 
the strength of the acid and the temperature, being prin- 
cipally antimony pentoxid, Sb 2 O s , when the acid is hot 
and concentrated. It is easily soluble in aqua-regia. 

In its physical properties antimony is a perfect metal, 
although it is a poor conductor of heat and electricity; 
but in its chemical properties it is only weakly metallic, 
so that in its compounds it generally acts as a non-metal, 
and closely resembles arsenic. 

The principal use of antimony is in the manufacture of 
alloys, the most important of which is type-metal. This 
is an alloy of lead, antimony, and tin, is very hard, and 
expands slightly when it solidifies, giving thus a perfect 
cast. It is also a constituent of the white or anti- friction 
metals (Compare 229). 

Antimony is used to some extent in the preparation of 
pharmaceutical compounds. 

COMPOUNDS OF ANTIMONY. 

404. Stibin, SbEL,. This substance, which is in every 



LECTURE NOTES 305 

way analogous to arsin (See 157), is formed when nascent 
hydrogen acts upon a soluble compound of antimony, thus: 

SbCl s + 6 H = SbH 3 + 3 HC1. 

Stibin is a colorless gas, having a very disagreeable 
odor and taste, and always contains some free hydrogen. 
It is somewhat soluble in water, and is poisonous, al- 
though not so much so as arsin. It is inflammable, and 
burns with a greenish-white flame, forming water, and 
antimony trioxid, Sb 2 O s . 

If stibin is made to pass through a glass tube heated 
to 150°, it is decomposed into antimony and hydrogen. 
The antimony is deposited on the sides of the tube, form- 
ing an antimony mirror. There are several reactions by 
which this may be distinguished from the arsenic mirror, 
which is formed in a similar way from arsin. The arsenic 
is easily volatilized, while the antimony does not volatil- 
ize, but fuses and forms small metallic globules; the arsen- 
ic is easily soluble in a solution of sodium hypochlorite, 
NaOCl, or of bleaching powder, while the antimony is 
insoluble in both (Compare 167). 

405. The Oxidsof Antimony. There are three oxids 
of antimony known. They are, 

Antimony trioxid, Sb 2 O s , 
Antimony tetroxid, Sb 2 4 , and 
Antimony pentoxid, Sb 2 6 . 

Antimony trioxid, Sb 2 O s , occurs in nature in the miner- 
als valentinite and se?iarmontite. It can be prepared by 
oxidizing antimony with dilute nitric acid ; also by dilut- 
ing a solution of antimony trichlorid, SbCl 3 , with hot 
water, washing the precipitate with sodium carbonate 
and igniting gently. 



306 LECTURE NOTES 

It is a white powder which is only very slightly soluble 
in water, and acts as a weak base. It dissolves in hydro- 
chloric acid, forming antimony trichlorid. It dissolves 
in tartaric acid, H 2 (C 4 H 4 6 ), and in acid potassium tar- 
trate, HK(C 4 H 4 6 ), forming potassium antimonyl tartrate, 
or tartar emetic, KSbO(C 4 H 4 6 ). When heated in the air, 
it absorbs oxygen, and forms antimony tetroxid. It 
volatilizes at a very high temperature and yields a vapor, 
the density of which corresponds to the molecular formula 
Sb ( 6 . 

Antimony tetroxid, Sb 2 4 , is regarded by some as a 
compound of antimony trioxid and pentoxid. It is formed 
when either of the two oxids are strongly heated in the 
air. It is a white powder which is very nearly insoluble 
in all acids. An impure form of this oxid occurs in nature 
as the mineral cervantite. 

Antimony pentoxid, Sb 2 O s , is formed by oxidizing anti- 
mony with hot, concentrated nitric acid, and heating the 
resulting metantimonic acid to 275°. It is a yellowish 
white powder, which is nearly insoluble in all acids. On 
strongly heating, it gives up a portion of its oxygen and 
forms the tetroxid. 

406. The Oxy-acids of Antimony. Although antimo- 
ny trioxid forms compounds in which it acts as a weak 
base, it also forms compounds with the strong alkaline 
bases in which it acts as a weak acid. When sodium 
hydroxid acts upon a solution of antimony, a white pre- 
cipitate is formed which is believed to have the composi- 
tion SbO(OH), or HSb0 2 . With an excess of sodium 
hydroxid, this precipitate dissolves, forming what is called 
sodium antimonite, NaSb0 2 . 

There are a few compounds known which are regarded 



LECTURE NOTES 307 

as derived from an acid corresponding to antimony tet- 
roxid. These are known as hypoa?ttimonates, and the 
potassium salt has the composition K 2 Sb 2 5 . 

There are three acids corresponding to antimony pent- 
oxid which are analogous to those of arsenic and phos- 
phorus. Ortho-antimonic acid, H 3 Sb0 4 , is obtained by de- 
composing potassium metantimonate with dilute nitric 
acid, and drying the compound at 100°. Pyro-antimonic 
acid is obtained by treating antimony pentachlorid with 
hot water and drying the precipitate at 100°, thus: 

2 SbCl 5 + 7 H 2 = H 4 Sb 2 7 + 10 HC1. 

When the pyro, or ortho-acids are heated to 200° they 
lose water and are converted into ?netantimonic acid, 
HSb0 3 . This acid is also formed when antimony is acted 
upon by hot, concentrated nitric acid. 

There are no salts of the ortho-acid known, and so the 
salts of the meta-acid are sometimes called antimonates. 
Acid sodium pyro-antimonate has the composition 
H 2 Na 2 Sb 2 O r It crystallizes with six molecules of water 
and is the least soluble of all the salts of sodium. 

407. The Halogen Compounds of Antimony. Antimo- 
ny forms compounds with each of the halogen elements, 
combining with them directly. The only ones of impor- 
tance are the chlorin compounds. 

Antimony trichlorid, SbCl 3 , is best formed by dissolving 
the metal, its trioxid, or trisulfid, in hydrochloric acid. 
It is a soft, deliquescent substance, which was known to 
the early alchemists, and called by them butter of antimo- 
ny. It melts at 73° and boils at 223°. It dissolves in a 
small quantity of water, but when added to a large quan- 
tity of water it forms a basic chlorid, SbOCl, which is 



308 LECTURE NOTES 

called antimony oxy-chlorid. This is a white, insoluble 
powder which was known to the alchemists as the powder 
of Algaroth. 

Antimony pentachlorid, SbCl 6 , is formed by the action 
of an excess of chlorin on antimony, or by leading chlorin 
into the fused trichlorid. It is a yellow, fuming liquid 
with a disagreeable odor. It forms a solid compound with 
a number of non-metallic chlorids and oxids, and is used 
in organic chemistry for chlorinating certain compounds. 

408. Antimony Trisulfid, Sb 2 S 3 . This is the most im- 
portant of all the compounds of antimony, and is found 
in nature as the mineral stibnite. It is frequently found 
in prismatic crystals, those from Japan being especially 
large, and fine. It has a dark-gray color and the crystals 
have a bright metallic luster. It is quite brittle, melts 
easily, and has a specific gravity of 4. 6. This substance 
has been known and used in the East since very early 
times under the name of Kohl. 

When hydrogen sulfid is passed into an acid solution of 
an antimony salt, an orange-red precipitate is formed. 
This is regarded as a hydrated form of antimony trisulfid; 
for, if this compound is heated in a current of carbon 
dioxid, it loses water, and passes into the dark-gray, 
crystalline variety. 

Antimony trisulfid was first employed as a medicine in 
the early part of the eighteenth century, by a Carthusian 
monk, and it was for a long time known as Carthusian 
powder. It was afterwards called kermes, and for many 
years was highly prized as a medicine. 

Antimony trisulfid is used in the manufacture of 
matches, and for making white fire. It is also used in 



LECTURE NOTES 309 

vulcanizing caoutchouc, to which it imparts a red color, 
and as a pigment in oil and water-color painting. 

409. Other Compounds of Antimony. When antimony 
trisulfid is dissolved in a solution of an alkaline sulfid, or 
is fused with an alkaline hydroxid, it forms the so-called 
sulfo-, or thio-antimonites (Compare 166). These are 
mostly brown or black compounds, which vary somewhat 
in composition, and are sometimes called the livers of an- 
timony. They decompose in hydrochloric acid, forming 
the trisulfid and sulfur. 

When hydrogen sulfid acts upon antimony pentachlorid, 
or any solution of the antimonates, it forms antimony 
pentasul/id, Sb 2 S 5 . This is an orange-red, rather unstable 
compound, which forms a series of sulfo- or thio-antimo- 
nates, analogous to those formed by the trisulfid. 

When antimony trioxid is boiled with acid potassium 
tartrate, there is formed potassium antimonyl tartrate, 
or tartar emetic KSbOCC^H^Og). This is a white crystalline 
compound which is easily soluble in water. It has been 
known since very early times, and used as a medicine, for 
which purpose it was highly valued by the iatro-chemists. 

Antimony compounds are all poisonous. They act upon 
the human system somewhat like arsenic, though not 
quite so violently. 

Bismuth. Symbol, Bi. Atomic Weight, 208.1. 

410. History and Occurrence. The writers of the thir- 
teenth century were the first to use the word marcasite, 
the name by which the metal bismuth was usually known 
until comparatively recent times; but as the name mar- 
casite was at first applied to almost any metallic looking 
ore, it is somewhat doubtful if bismuth was known to the 



310 LECTURE NOTES 

thirteenth century writers. It was, however, clearly 
known to the later alchemists, probably as early as the 
seventeenth century. 

Bismuth is a comparatively rare metal. It is usually 
found free in nature, but occurs also in a number of min- 
erals, the most important of which are bismuthinite, Bi 2 S 3 , 
and bismite, Bi 2 O r Most of the bismuth comes from the 
native ore, which is found principally in Saxony, in Peru 
and Bolivia, and in Australia. 

411. Preparation, Properties, and Uses. In the prepa- 
ration of bismuth from its ores, these are first roasted, 
and then smelted with iron, carbon, and some slag, after 
which the easily fusible metal is drawn off. This crude 
metal, or the native ore, is further purified by heating on 
an inclined iron plate when the bismuth melts and runs 
away very nearly pure. Since bismuth is employed for 
making pharmaceutical preparations, it is often necessary 
to have it very pure. To obtain this, the crude metal is 
dissolved in nitric acid, forming bismuth nitrate, Bi(NO s ) s 
and a basic nitrate, Bi(OH) 2 N0 3 , is then precipitated by the 
addition of water. The basic nitrate is dried, changed to 
the oxid by heating, and reduced to the metal with carbon. 
If the metal contains arsenic, this may be removed by 
fusing with potassium nitrate. 

Bismuth is a grayish-white metal, with a distinctly 
reddish tinge. It is very hard and brittle, and has a 
bright metallic luster. Its specific gravity is 9.8 and it 
melts at 265°, the fused metal expanding slightly on 
solidifying. If the fused metal is cooled until it begins 
to solidify, and the remaining liquid is then poured off, 
fine rhombohedral crystals are obtained, which closely 
resemble cubes. 



LECTtJRE NOTES 31 1 

Bismuth remains unchanged in dry air at the ordinary 
temperature, but in moist air its surface becomes tar- 
nished, and, when strongly heated, it burns, forming 
bismuth oxid, Bi 2 O s . It is insoluble in hydrochloric acid, 
but dissolves in hot concentrated sulfuric acid. It is 
easily soluble in cold nitric acid, and in aqua regia. 

Bismuth is principally employed in making the so-called 
fusible alloys such as Newton's metal, Wood's metal, and 
others (See 228). Some of its compounds are used in 
medicine. 

Bismuth forms no compound with hydrogen. In other 
respects its compounds resemble those of antimony. 

COMPOUNDS OF BISMUTH 

412. The Oxids of Bismuth. Bismuth forms four com- 
pounds with oxygen. These are : 

Bismuth dioxid, Bi 2 O a , 
Bismuth trioxid, Bi 2 O g , 
Bismuth tetroxid, Bi 2 4 , and 
Bismuth pentoxid, Bi 2 5 . 

None of these compounds is an acid-forming oxid, and 
they all form the same series of salts when acted upon by 
acids, the bismuth acting always as a trivalent metal. 

Bismuth trioxid, Bi 2 O s , which is the common oxid, is 
the most stable and important of the four. When the 
others are heated in the air, they each form the trioxid. 
It is found in nature as the mineral bismite, and is formed 
artificially when the metal, the nitrate, or the carbonate, 
is strongly heated in the air. 

Bismuth trioxid is a light yellow, amorphous powder, 
which is insoluble in water, but dissolves in acids, form- 
ing salts of bismuth. 



312 LECTURE NOTES 

413. Bismuth Nitrate, Bi(N0 3 ) 3 . This compound is 
formed when the metal, or its oxid, is dissolved in nitric 
acid. It forms large tabular crystals, containing five 
molecules of water. It is soluble in a small quantity of 
water, but when an excess of water is added a basic ni- 
trate is precipitated. 

Basic bismuth nitrate, Bi(OH) 2 NO s , which is known 
commercially as sub?iitrate of bismuth, was known to the 
alchemists, and highly valued by them as a medicine. It 
is still largely used for this purpose in cases of chronic 
diarrhoea, and cholera, and is also used as a cosmetic. It 
is also used for producing an iridescent glaze on glass or 
porcelain, and for staining glass. 

414. Other Compounds of Bismuth. Bismuth forms 
compounds with the halogen, and with most of the other 
acids. They are all unstable, and decompose on the addi- 
tion of water, forming insoluble basic compounds. 

Bismuth trichloride BiCl 8 is formed when bismuth is 
heated in a stream of chlorin, or when bismuth trioxid is 
dissolved in hydrochloric acid. It is a white crystalline 
compound, which is quite deliquescent. Water decom- 
poses it forming bismuth oxychlorid, BiOCl. 

Bismuth sulfid, Bi 2 S 8 , is found in nature as the mineral 
bismuthinite. It is prepared artificially by fusing bismuth 
with an excess of sulfur, or by leading hydrogen sulfid 
into a solution of a bismuth salt. It is a dark-brown 
amorphous compound which is insoluble in the alkaline 
hydroxids and sulfids (Compare 409). 

The other compounds are not important. 



LECTURE NOTES 313 

THE METALS OF GROUP VI. 

415. General Characteristics. The two divisions which 
make up this group are very well denned. Division A 
consists of the elements chromium, molybdenum, neody- 
mium (?), tungsten and uranium. With the exception of 
neodymium, about which very little is known, the ele- 
ments of this division are all well denned metals. Chro- 
mium is not a common substance, and the other metals are 
all quite rare. They do not combine with hydrogen, but 
each forms several oxids, one of which is quite character- 
istic and corresponds to the general formula MO s , the 
metal being sexivalent. It is also an acid-forming oxid, 
and combines with the other metals to form a well defined 
series of salts. 

Division B consists of the elements oxygen, sulfur, se- 
lenium and tellurium. These are all well defined non- 
metals and have already been described. The valence of 
the elements of this group varies from two to six. 

Chromium. Symbol, Cr. Atomic Weight, 52.1. 

416. History and Occurrence. This metal was discov- 
ered in the mineral crocoite in 1797, by Vauquelin. He 
found that the mineral contained lead, combined with an 
unknown acid, which he recognized as the oxid of a new 
element. To the new element he gave the name chromium, 
because all of its compounds were colored. During the 
same year, the new metal was discovered independently 
by Klaproth. 

Chromium never occurs free in nature. It occurs most 
abundantly in the mineral chromite, FeO, Cr 2 O s , from 
which most of the chromium and its compounds are ob- 
tained. It is also found in the mineral crocoite, PbCrO,, 



314 LECTURE NOTES 

and in a few other rare minerals, and it is the coloring ma- 
terial in a number of minerals, such as beryl {emerald), 
serpentine, penninite, and others. 

417. Preparation, Properties, and Uses. Chromium 
was first obtained by strongly heating the oxid with car- 
bon, in a lime crucible. It may also be obtained by heat- 
ing the chlorid with zinc, under a layer of sodium chlorid, 
and dissolving out the zinc with nitric acid. It is best 
obtained by heating the oxid with carbon in an electric 
furnace, or, by reducing the oxid with powdered aluminum 
(Compare 355). 

Chromium is a very hard, steel-gray, crystalline metal, 
which much resembles iron, but is somewhat whiter. It 
has a specific gravity of 6.9, and melts at a temperature 
much higher than is required to melt platinum (about 
2000°). It is permanent in dry air at all ordinary tem- 
peratures, but will burn in the oxyhydrogen flame. It 
dissolves slowly in dilute mineral acids, but is hardly 
affected by hot, concentrated nitric acid. 

Chromium has little technical importance. "About its 
only use is in the manufacture of chromium steel, which 
is made by the addition of from two to four per cent, of 
chromium to ordinary steel. It is exceedingly hard and 
tenacious, and is used in making the shells for projectiles. 

COMPOUNDS OF CHROMIUM. 

418. The Oxids of Chromium. There are two, well 
defined oxids of chromium known. These are: 

Chromic oxid, or Chromium sesquioxid, Cr 2 O s , and 
Chromium trioxid, CrO s . 

The former is a basic oxid, and forms with all of the 



LECTURE NOTES 315 

stronger acids a series of salts in which the chromium is 
trivalent. The latter is an acid-forming oxid, and com- 
bines with water to form chromic acid, H 2 Cr0 4 , and with 
bases to form a series of salts called chromates, in which 
the chromium is sexivalent. 

In addition to these two oxids there exists an hydroxid, 
Cr(OH) 2 , corresponding to an oxid, CrO, and a very un- 
stable acid, called perchromic acid, HCr0 4 , corresponding 
to an oxid, 0r 2 O 7 ; but neither oxid exists alone. 

The two known oxids, together with the two hypothet- 
ical oxids, are each the basis for a series of chromium 
compounds, in two of which the chromium is basic, while 
in the other two it is acid. 

The salts corresponding to the hypothetical oxid, CrO, 
are called chromous salts. They are very unstable, and. 
although several of them are known, they possess only 
theoretical importance. 

The salts corresponding to the two known oxids are 
called chromic salts (more commonly chromium salts), and 
chromates, respectively. Both series of compounds are 
quite stable, and a number of the salts are quite impor- 
tant. 

The salts corresponding to the other hypothetical oxid, 
Cr 2 7 , are called per chromates. These are very unstable, 
only a single sodium salt having as yet been prepared, 
and that of somewhat doubtful composition. 

419. Chromic Oxid, Cr 2 3 . This compound may be 
prepared by heating chromium hydroxid, Cr(OH) 3 , and 
also by heating ammonium dichromate, thus: 

(NH 4 ) 2 Cr a O, = Cr 2 O s + 4H.O + N,, 



316 LECTURE NOTES 

It is a green, amorphous powder, which becomes crys- 
talline after fusing in the oxyhydrogen flame. The crys- 
talline form is insoluble in acids. 

Chromic oxid is used for coloring glass and porcelain, 
and imparts to them a fine green color. It is also used 
as a pigment in ordinary painting, and is known as chrome- 
green. 

There are several hydroxids corresponding to this oxid. 
When an alkaline hydroxid is added to a solution of a 
chromium salt, the normal chromium hydroxid, Cr(OH) s , 
is formed. When this compound is heated to 200° in a 
stream of hydrogen, it loses water and forms a compound, 
CrO(OH), which acts like a weak acid, and forms a series 
of compounds with the stronger bases. These are called 
chromites, and correspond to the aluminates (See 358). 

When potassium dichromate is fused with boric acid, 
there is formed a beautiful green compound, having the 
composition Cr 2 0(OH) 4 . This always contains some boric 
acid, and is used as a pigment under the name of Guig- 
netfs green. 

420. Chromium Trioxid, Cr0 3 . When a solution of a 
chromate is acted upon by concentrated sulfuric acid, it 
is decomposed, and chromium trioxid is formed, thus: 

K 2 Cr0 4 + H 2 S0 4 = K 2 S0 4 + H,0 + CrO s . 

Chromium trioxid forms long, red needles or prisms, 
which are easily soluble in water. When heated in the 
air to about 250° it decomposes into chromic oxid and 
oxygen. It is, therefore, a powerful oxidizing agent, and 
is much used in organic chemistry for this purpose. 

When chromium trioxid is warmed with too little water 



LECTURE NOTES 317 

to effect complete solution, it forms chromic acid, H 2 CrO^ 
On cooling the solution, the acid separates out in rose 
colored crystals, which are very unstable. It is believed 
that with excess of water the solution of this oxid exists 
in the form of dichromic acid, H 2 Cr 2 7 . Each of these 
acids forms a series of salts, called chromates, and dichro- 
mates, respectively, some of which are of considerable 
importance. 

421. Potassium Dichromate, K 2 Cr 2 7 . This most im- 
portant compound is known commercially as bichromate 
of potash, and is used for the preparation of nearly all 
other chromium compounds. It is made on a large scale 
from the mineral chromite, called also chrome-iron ore. 
by roasting the finely crushed ore mixed with potassium 
carbonate and lime, in a reverberatory furnace, the action 
being as follows: 

2 FeO, Cr 2 O s -f 4 K 2 C0 3 + 4 CaO -f 7 O = 

4 K 2 Cr0 4 + 4 CaC0 3 + Fe 2 3 . 

The yellow mass is then lixiviated with water, and 
mixed with the proper amount of sulfuric acid, when the 
normal potassium chromate is changed to the dichromate. 
thus: 

2 K,00 4 + H^SO, = K 2 0r 2 O, + K,SO, + H 2 0. 

From this solution the potassium dichromate crystal- 
lizes in fine garnet-red crystals, which contain no water 
of crystallization. It is easily soluble in water and the 
solution gives an acid reaction. It has a bitter, metallic- 
taste, and is quite poisonous. It fuses without decomposi- 
tion at a high temperature, but at a white heat it decom- 
poses giving off oxygen. It oxidizes organic matter, 
especially when exposed to the light, and at the same 



318 LECTURE NOTES 

time is itself reduced to the oxid. This action is utilized 
in photography in the so-called Autotype process. 

Potassium dichromate is used for the preparation of 
various chromates some of which have a bright color and 
are used as pigments (Compare 393). 

422. Chromium Sulfate, Cr 2 (S0 4 ) 3 . This compound is 
formed by dissolving chromium hydroxid in sulfuric acid. 
The solution is at first green, but after standing for a 
time it becomes violet, and the chromium sulfate is depos- 
ited as a violet, crystalline mass, containing eighteen 
molecules of water of crystallization. If the violet solution 
is warmed for a time it again becomes green, but after 
cooling and standing, it again changes to violet. Recent 
investigation of this green compound seems to show that 
the chromium, and a part of the sulfuric acid radical, ex- 
ist as a complex basic radical called sulfo-chromyl, of 
which the green compound is the sulfate, its composition 
being [Cr 0(SO ) ] "S0 4 . Other chromium salts show an 
analogous action. 

Chromium sulfate combines with the alkaline sulfates 
to form alums, the potassium compound being called 
chrome- a I urn (Compare 361). 

Molybdenum. Symbol, Mo. Atomic Weight, 96. 

423. Molybdenum and its Compounds. This element 
was first obtained in 1782, by Hjelm. It does not occur 
free in nature, and its compounds are quite rare. It oc- 
curs principally as the minerals moly bdenite, MoS 2 , and 
wulfenite, PbMo0 4 . 

The metal can be prepared by igniting the chlorids or 
oxids in a stream of hydrogen; also by igniting the oxids 
with carbon in an electric furnace. It is a hard, mallea- 



LECTURE NOTES 319 

b le metal, having a specific gravity of about 9, and fuses 
at a higher temperature than platinum. 

Molybdenum is interesting from the great variety of 
its compounds. It forms four compounds with chlorin, 
viz. : MoCl 2 , MoCl g , MoCl 4 , and MoCl 5 . It also forms four 
compounds with oxygen, viz. : MoO, Mo 2 O s , Mo0 2 . and 
Mo0 3 . The first three are basic oxids, while the fourth 
is an acid-forming oxid, combining readily with the bases 
to form salts which are called molybdates, and which are 
analogous to the chromates. 

Molybdenum trioxid, Mo0 3 , is formed by roasting molyb- 
denite in the air. It is a white amorphous compound 
which is insoluble in water and in acids, but dissolves 
readily in the alkalies and in ammonium hydroxid. 

Normal ammonium molybdate, (NH 4 ) 2 Mo0 4 , is form- 
ed by dissolving the trioxid in concentrated ammo- 
nium hydroxid in excess. When this solution is evapo- 
rated, large colorless crystals are formed, which have the 
composition, (NH 4 ) 6 Mo 7 24 , 4 B 2 0. and which are com- 
monly called ammonium molybdate. 

When a solution of ammonium molybdate in an excess 
of nitric acid, is added in excess to a solution containing 
phosphoric acid, or a phosphate, a yellow compound is 
formed. This is known as ammonium p ho sp ho -molybdate. 
and has the composition, (NH 4 ) 3 P0 4 , 12 Mo0 3 . A similar 
compound is formed with arsenic acid. Some of the molyb- 
dates are very complicated. 

Tungsten. Symbol, W (Wolframium). Atomic Weight, 
189. 

424. Tungsten and its Compounds. This rare element 
was first obtained in 1783, by the Spanish chemists 



320 LECTURE NOTES 

d'Elhujar. It occurs principally in nature in the form 
of tungstates, the more important ones being the miner- 
als wolframite, (FeMn)W0 4 , scheelite, CaW0 4 , and stolzite, 
PbW0 4 . 

The metal is obtained by methods exactly analogous to 
those given under molybdenum, which metal tungsten very 
closely resembles. It is a very hard, steel-gray metal, 
having a specific gravity of 19.1, and fuses at a higher 
temperature than molybdenum. The metal is used for the 
manufacture of tungsten-steel, an alloy which contains from 
eight to ten per cent, of tungsten, and is exceedingly 
hard. 

Tungsten forms but two oxids, W0 2 and W0 3 , the latter 
being an acid-forming oxid, forming salts which are called 
tungstates, and which are analogous to the molybdates. 

The partial reduction of the alkaline tungstates results 
in a peculiar class of compounds, which have various colors 
and a metallic luster, and are called the tungsten bronzes. 

Uranium. Symbol, U. Atomic Weight, 239.6. 

425. Uranium and its Compounds. This element was 
first obtained in 1842, by Peligot. It is not a very abun- 
dant element, and occurs in nature chiefly as uraninite, or 
pitch- blende, U 3 O g , and in small quantities in a few rare 
minerals. The uranium minerals nearly always contain 
the rare element helium. 

Uranium can be prepared by the action of sodium or 
potassium on uranium tetrachlorid. It is best prepared 
by reducing the oxid with carbon in an electric furnace. 

It is a white metal, and has a specific gravity of 18.7. 
When in a powdered form, it oxidizes easily in the air, 



LECTURE NOTES 321 

and decomposes water. It melts at about 1500°, and 
volatilizes at a very high temperature. 

Uranium forms two oxids, U0 2 , and U0 3 , and these 
combine to form other oxids. The dioxid is a basic oxid, 
and forms a series of uranous salts, in which the metal is 
quadrivalent. The trioxid is an acid-forming oxid and 
forms the salts called uranates, and also acts like a basic 
oxid in forming a series of salts, in which one of the oxy- 
gen atoms is replaced by an acid radical. These are also 
regarded as salts of the basic radical uranyl, (U0 2 )". 

The compounds of uranium all have a color which is 
usually a shade of green or yellow, and most of them show 
a strong fluorescence. Glass colored by uranium oxid 
is known as uranium glass, and has a beautiful greenish- 
yellow color, which appears yellow by transmitted light, 
and green by reflected light. 

Uranium has the highest atomic weight of any known 
element. 

THE METALS OF GROUP VII. 

426. General Characteristics. The two divisions of the 
seventh group are very distinct in their properties. Di- 
vision B consists of the four very characteristic non-met- 
als, fluorin, chlorin, bromin, and iodin, which are known 
as the halogen elements, and which have already been 
described. 

Division A consists of the metals manganese, an undis- 
covered element having an atomic weight of about 100, 
and possibly samarium. The latter element has been 
placed in this group in the table on page 192, because, 
from its atomic weight, it would seem to belong here; but 
its position in the periodic system is uncertain, and too 



322 LECTURE NOTES 

little is known of it to make it certain that it is really an 
element and not a mixture of elements. Manganese is at 
present, therefore, the only representative of the metals 
in this group. 

The relationship of the two divisions of this group is 
best seen in their highest state of oxidation. Permanga- 
nic acid, HMnO^, is exactly analogous to perchloric and 
periodic acids, and the salts of these three acids are quite 
similar, and for the most isomorphous. The principal 
element in each of these acids appears to be septivalent. 

Manganese. Symbol, Mn. Atomic Weight, 55. 

427. History and Occurrence. The compound manga- 
nese dioxid, Mn0 2 , which is found in nature as the min- 
eral pyrolusite, has been known since very early times. 
It was for a long time believed to be a variety of magnet- 
ite, or loadstone, which mineral it somewhat resembles, 
and it is probably for this reason that it was described by 
Pliny as one of several kinds of magnes, or magnets, but 
which, as he says, " being of the feminine gender, does 
not attract iron." The metal was first obtained by Gahn, 
in the latter part of the eighteenth century. 

Manganese never occurs free in nature, but its com- 
pounds are quite widely distributed, and are principally 
oxids. In addition to pyrolusite the most important min- 
erals are manganite, MnO(OH), braunite, Mn 2 O s , haus- 
mannite, Mn 3 4 , rhodonite, MnSiO g , and rhodocrosite, 
MnCO g . It is also found as the coloring matter in a num- 
ber of minerals, the best known being the variety of 
quartz called amethyst. 

428. Preparation, Properties, and Uses. Manganese 
is usually prepared by reducing the oxids by means of 



LECTURE NOTES 323 

carbon. The process is a difficult one, and requires a 
very high temperature, which is best obtained in an elec- 
tric furnace. The metal thus obtained always contains 
considerable carbon, and is called cast -manganese. A 
purer form may be obtained by reducing the oxid with 
aluminum. 

Manganese is a hard, gray metal, with a reddish tinge. 
It has a specific gravity of about 8, and melts at about 
1900°. The pure metal oxidizes easily in moist air, and 
decomposes water, even in the cold. It dissolves easily 
in the common acids, and forms salts in which the metal 
is bivalent. Cast-manganese usually contains some iron, 
and is permanent in the air. 

It is used in the manufacture of steel, manganese steel 
being much harder and less easily oxidized than ordinary 
steel. An alloy of manganese with iron and copper is 
called manganese bronze, and is used for making the pro- 
pellers of steamers. 

COMPOUNDS OF MANGANESE. 

429. The Oxids of Manganese, Manganese forms six 
compounds with oxygen. These are: 

Manganese monoxid, or Manganous oxid, MnO, 
Mangano-manganic oxid, or Red manganese oxid, Mn 3 4 , 
Manganese sesquioxid, or Manganic oxid, Mn 2 O s , 
Manganese dioxid, Mn0 2 , 
Manganese trioxid, Mn0 3 , and 
Manganese heptoxid, Mn 2 O r 

The monoxid and sesquioxid are basic oxids, and give 
rise to two series of salts called manganous and manganic 
salts, respectively. The red oxid, Mn 3 4 , is also a basic 



324 LECTURE NOTES 

oxid, and yields both kinds of salts, it being generally re- 
garded as a combination of the monoxid and sesquioxid, 
MnO,Mn 2 O s , or of the monoxid and dioxid, 2MnO,Mn0 2 . 

The dioxid forms a series of salts with the strong bases 
which are called manganites, while with acids it forms 
manganous salts with elimination of oxygen. The trioxid 
and heptoxid are both acid -forming oxids, and form salts 
which are called manganates and permanganates, respec- 
tively. 

When a soluble hydroxid is added to a manganese solu- 
tion, manganous hydroxid, Mn(OH) 2 , is formed. This 
oxidizes easily in the air, and forms manganic hydroxid, 
MnO(OH). 

430. Manganese Dioxid, Mn0 2 . This compound which 
is also called manganese peroxid, and is known commer- 
cially as the black oxid of manganese, is found in nature 
as the mineral pyrolusite. It is a hard, black, crystalline 
solid, and is the most important ore of manganese. This 
compound can be formed artificially by heating manga- 
nous nitrate, Mn(N0 3 ) 2 , or carbonate, MnCO s , in the air, 
and by several other methods. The artificial product is 
a dark brown or black amorphous powder. 

When manganese dioxid is strongly heated, it decom- 
poses into mangano-manganic oxid, Mn 3 4 , and oxygen. 
It is the principal source of manganese compounds and is 
used in the manufacture of colorless glass. The most 
important use of this compound is for making the chlorin 
which is used in the manufacture of bleaching powder. 

431. The Weldon Process. In making chlorin by the 
action of hydrochloric acid on manganese dioxid. the 
manganese which remains in the solution in the form of 



LECTURE NOTES 325 

manganous chlorid (See 64), was at first allowed to run 
to waste; but the relatively high cost of the manganese 
dioxid has resulted in the invention of a number of pro- 
cesses for recovering the waste manganese. The best of 
these was invented by Weldon, in 1867. 

By this process, the waste liquid is first mixed with 
ground .chalk, or limestone, which neutralizes the excess 
of hydrochloric acid. The liquid is then allowed to settle, 
after which the clear liquid is pumped into an open tank, 
called the oxidizer , where it is mixed with an excess of 
milk of lime, Ca(OH) 2 . This precipitates the manganese 
as hydroxid, thus: 

MnCl 2 -f Ca(OH) 2 == Mn(OH) 2 + CaCl 2 . 

The manganous hydroxid, with the excess of lime is 
then heated to a temperature of 55°, and left exposed to 
the air for some time, when it oxidizes, forming calcium 
manganite, thus: 

Mn(OH) 2 + Ca(OH) 2 -f O = CaMn0 3 + 2 H 2 0. 

If the milk of lime is not in sufficient excess, other 
manganites are formed, such as CaMn 2 O s and CaMn 5 O n . 

The contents of the oxidizer are then run out into a 
series of tanks called mud settlers, when the calcium man- 
ganite settles as a thin black mud, known as Weldon mud. 
This compound, which may be regarded as a mixture of 
lime and manganese dioxid, CaO -j- (Mn0 2 ) x gives chlorin 
when acted upon by hydrochloric acid. 

432. The Manganates. The salts known as manga- 
nates, are derived from the hypothetical manganic acid. 
H 2 Mn0 4 . Although this acid has never been obtained free 
its anhydrid is known. This is called manganese trioxid, 



32(3 LECTURE NOTES 

Mn0 3 , and is a reddish, amorphous compound. It is very 
unstable, and decomposes when dissolved in water, form- 
ing permanganic acid, HMn0 4 , thus: 

3 MnG 3 + H 2 = 2 HMn0 4 + Mn0 2 . 

The alkaline manganates can be obtained by fusing 
manganese dioxid with an alkaline hydroxid, or carbonate, 
in the presence of some oxidizing agent, thus: 

3 MnO, 4- 6 KOH 4- KCIO, = 3 KMnO, 4- KC1 4- 3 HO. 

The fused mass has a dark-green color, and dissolves in 
a small amount of cold water forming a green solution, 
which is only stable in the presence of some free alkali. 
When the manganates are dissolved in considerable water, 
or in an acid solution, they are decomposed forming per- 
manganates, the color changing from green to red. 

433. The Permanganates. The salts known as per- 
manganates are derived from permanganic acid, HMn0 4 . 
If potassium permanganate is added to well cooled con- 
centrated sulfuric acid, olive-green, oily drops of manga- 
nese heptoxid, Mn 2 7 , are obtained. These dissolve in 
water, forming a purple solution of permanganic acid. 
Both the oxid and the acid are very unstable, and, when 
heated, decompose with explosive violence, forming man- 
ganese dioxid, and oxygen. The permanganates, especi- 
ally those of the alkalies, are quite stable, and are of 
considerable importance. 

Potassium permanganate, KMnO^, is the most impor- 
tant of this class of salts. It is prepared by oxidizing 
manganese dioxid in the presence of some strong base. 
Usually potassium chlorate is the oxidizing agent and 
potassium hydroxid is the base. [The equation for this 



LECTURE NOTES 327 

reaction is given in the preceding article.] The manganate 
thus formed is decomposed by dissolving the mass in wa- 
ter, and passing carbon dioxid through the solution, thus: 

3 K 2 Mn0 4 + 2 C0 2 = 2 KMn0 4 + Mn0 2 + 2 K 2 C0 3 . 

Potassium permanganate is a dark, crystalline sub- 
stance, which at first has a greenish, metallic luster, but 
which becomes steel-blue on exposure to the air, after 
which the salt suffers no further alteration. The crystals 
have a specific gravity of 2.7, arid, when powdered, have 
a reddish-purple color. The salt is used extensively as 
an oxidizing agent. In the presence of an acid, two mol- 
ecules of the salt yield five atoms of oxygen to the com- 
pound which is being oxidized, the manganese forming a 
manganous salt, thus: 

2 KMn0 4 + 3 H 2 S0 4 = 2 MnS0 4 -f K 2 S0 4 + 3 H 2 + 5 O. 

In the presence of an alkali, the salt is reduced to a 
manganite, two molecules of the salt yielding three atoms 
of available oxygen, thus : 

2 KMn0 4 + 2 KOH = 2 K 2 Mn0 3 + H 2 + 3 0. 

A solution of crude sodium permanganate is quite ex- 
tensively employed for disinfecting purposes, under the 
name of Condy's 



434. The Manganous Salts. These salts are generally 
quite stable, but they have little or no commercial value. 
They are formed by dissolving the metal, its oxids, or car- 
bonate, in the different acids. The salts are usually pale- 
red, or white, and the manganese in them is bivalent. 

Manganous chlorid, MnCl 2 , forms rose colored crystals 
with four molecules of water. It forms double salts with 
the alkaline chlorids. 



;_ - ::::":: rz 



: : _ : -r - : ■: : i '. - 



--1- 



MoCO r occur- i ■ 



tI--..": " 



:: - - : i " i - 



lectuee yorz- 329 

THE METALS OF GROUP VIII. 

436. General Characteristics. In arranging the ele- 
ments according to the periodic system as given on page 
192, three metals are found in the midst of the first long 
period, between manganese and copper, and connecting 
the fourth and fifth series. In like manner three elements 
are found connecting the sixth and seventh series, and 
three more connecting the tenth and eleventh series, that 
is. in the midst of the second and fourth long periods. 
To make the system complete there should be three ele- 
ments connecting the eighth and ninth series in the third 
long period, and it seems quite probable that such ele- 
ments exist: but the elements which belong to this perk 
are many of them rare earths (351), about which compar- 
atively little is known, and at present no connecting 
elements are known. 

It has already been noted _ .bat som^ rfore 

the discovery of the rare gases, recently found in the at- 
mosphere, it was predicted that elements existed which 
would be found to connect the very strongly negative 
halogen elements of Group VII. and the very strongly 
positive alkali metals of Group I. 

The eighth group is. therefore, a group of connecting 
elements. It classifies itself naturally into two divi- 
one. which we will call division A. consists of the elements 
which connect each period with the following : the other, 
division B. consists of the three sub-divisions, of three 
elements each, which connect the two series of the long 
periods. 

Division A consists of the elements helium, neon, ar^ n 
krypton, and xenon. These are all extremelv rare ele- 



330 LECTURE NOTES 

ments, and are distinguished by their entire lack of 
valence, that is, they cannot be made to combine with 
anything. They are non-metallic, and all of them are 
gases, but too little is known about them at present, to 
warrant more than the mention of them at this point. 

Division B divides itself naturally into the sub-divisions 
of three elements which occur in the midst of each period. 
The first sub-division is called the iron group, and consists 
of the metals iron, nickel and cobalt. They are much 
alike in their physical properties, and are fairly common. 
The second sub-division is called the ruthenium group, or 
the light platinum metals, and consists of the metals ru- 
thenium, rhodium, and palladium. These are very rare 
metals, and are usually found associated with the metals 
of the third sub-division. The latter is called the plati- 
num group, or the heavy platinum metals, and consists of 
the rare metals osmium, iridium, and platinum. 

The homologous elements, iron, ruthenium, and osmium, 
show similarity in their chemical properties, their highest 
oxids being acid-forming. This property diminishes in 
the other members of each sub-division, and the third 
series of homologous elements is much like the metals of 
division B of Group I. 

437. The Iron Group. The three metals of this sub- 
division are distinguished from all other elements by be- 
ing magnetic. They are all oxidized when strongly heated 
in the air, or oxygen, and at a red heat decompose water. 
They all form two series of compounds, corresponding to 
the basic oxids M"0, and M'" 2 O s , and, in addition, iron 
forms a series of salts called ferrates, corresponding to an 
hypothetical acid-forming oxid FeO g . Both ferrous and 
ferric compounds are quite stable, the ferric compounds 



LECTURE NOTES 331 

being the more stable. Nickel and cobalt being more 
basic, their lower compounds are much the more stable, 
their higher compounds being so unstable, that, for the 
most part, they have only a theoretical importance. 

The metals of Group VIII show a tendency to form 
complex radicals. This is seen in this sub- division in the 
cyanogen compounds of iron, and in the cobaltamines. 

Iron. Symbol, Fe (Ferrum). Atomic Weight, 56. 

438. History and Occurrence. Iron is by far the most 
important of all the metals. It has been known since 
very early times. It is mentioned in several places in the 
Bible, and was known to the ancient Assyrians and 
Egyptians. The Greeks and Eomans also knew it, and 
some of the mines from which they obtained their ore are 
still in operation. The early alchemists considered iron as 
sacred to Mars, the god of war. They connected it also 
with the planet Mars, and gave to it the symbol of the 
planet, ^, which is supposed to represent a spear and 
shield. 

Iron is the fourth most widely distributed element 
known, and constitutes about 5.5 per cent, of the earth's 
crust. It is occasionally found free in nature. The na- 
tive iron occurs principally in meteorites, where it is 
always alloyed with more or less nickel, and generally 
some other metals. It is also found in small particles in 
the Missouri coal measures, and also in certain igneous 
rocks. 

Iron is usually found in combination with oxygen or 
sulfur, and occasionally with other elements. It is also 
found in small quantities in almost every known rock and 
mineral. 



332 LECTURE NOTES 

Among the great number of iron minerals the following 
maybe mentioned as most important: magnetite. Fe s 4 ; 
hematite, Fe 2 O s ; limonite, Fe 4 3 (OH) g ; pyrite, FeS 2 ; and 
siderite, FeC0 3 . 

Iron is also found in the blood and in the ashes of plants. 

439. Preparation. In describing the preparation of 
iron, we have to distinguish between pure iron, and the 
iron which is used for commercial purposes. Absolutely 
chemically pure iron is almost unknown, but can be ob- 
tained by reducing pure iron oxid, or chlorid, in a stream 
of hydrogen, and also by electrolysis. 

Commercial iron is obtained by reducing the oxid by 
means of carbon, in a blast furnace. This form of iron 
always contains some impurities, and especially carbon, 
of which it may contain as high as five per cent. The 
carbon contained may be chemically combined, as iron 
carbid, Fe s C, or mechanically mixed, as graphite (174), or 
both. The amount of carbon present, and the form in 
which it exists, together with the presence of certain 
metals as impurity, results in the varieties called cast 
iron, wrought iron, and steel, with their various modifi- 
cations. 

440. Properties. Pure iron is a hard, grayish-white 
metal, which is very ductile and tenacious. It has a spe- 
cific gravity of 7.8, and melts at about 2000°. At a red 
heat it becomes soft, and at a white heat it can be welded. 
Pure iron is attracted by a magnet; and if an electric 
current is made to pass through a wire wound around a 
bar of iron, the iron becomes a magnet, but loses this 
power as soon as the current is stopped. 

Iron remains permanent in dry air but in moist air it 



LECTURE NOTES 333 

quickly oxidizes, or rusts. In order to avoid this, iron 
implements are often coated with oil, fat, or graphite, or 
painted with oil paints. If iron is exposed to steam, 
heated to a temperature of 650°, it becomes coated with 
a film of black magnetic iron oxid, Fe 3 4 which prevents 
further oxidation. For many purposes, iron is covered 
with zinc, tin, or nickel, to prevent rusting. 

Iron combines directly with the halogens, and, at a red 
heat, with sulfur and oxygen. It also decomposes water 
at a red heat. It dissolves easily in most dilute acids, 
forming ferrous salts, and hydrogen. With dilute nitric 
acid, it forms ferrous nitrate, and ammonium nitrate, 
thus: 

4 Fe + 10 HNO, = 4 Fe(N0 3 ) 2 + NH 4 N0 3 +3 H 2 0. 

With concentrated nitric acid it forms ferric nitrate, 
Fe(NO s ) 3 , and nitric oxid. 

When iron is dipped into concentrated nitric acid, and 
then washed in water and dried, it becomes passive, and 
is no longer soluble in nitric acid. This is probably due 
to the formation of a thin film of oxid on the surface. 

441. The Metallurgy of Iron. The metallurgy of iron 
forms one of the most important industries of modern 
times. The reduction of the ores is carried on in what is 
called a blast furnace. This is a tall shaft, from 50 to 100 
feet high and from 15 to 25 feet in diameter in its widest 
place. At the bottom of the furnace is the hearth or 
crucible, a cylinder from 8 to 12 feet in diameter and about 
the same in height. On this rests the shaft or chimney, 
which is, in form, two truncated cones joined at their bases. 
The upper cone is called the "body "and the lower one 
the " boshes." The shaft is usually constructed with an 



334 LECTURE NOTES 

outer shell of wrought iron, supported on pillars, and 
lined with fire-bricks, the whole being encased in a wall of 
common bricks. 

The blast is introduced into the furnace through the 
walls of the hearth, by means of tubes called tuyeres. 
These are made of wrought iron, and vary in number ac- 
cording to the size of the furnace. Near the bottom of 
the hearth is an opening, where the molten iron is drawn 
off, and above is another opening, through which the slag 
is allowed to flow away. 

Before introducing the iron ore into the furnace, it is 
broken up and roasted, in order to drive off water, to de- 
compose carbonates, to oxidize sulfids, and, in general, to 
change the ore as much as possible, into ferric oxid,Fe 2 3 , 
which is the most easily reduced of any of the oxids of 
iron. 

Considerable care has to be exercised in charging the 
furnace for the first time, and it often takes months 
before the furnace is doing full work. A fire is built on 
the hearth, and maintained until the walls of the furnace 
are thoroughly heated. The furnace is then filled with 
alternate layers of carbon, in the form of coke or charcoal, 
and the ore, mixed with the proper fluxes. The fluxes 
depend upon the nature of the ore. If, as is often the 
case, the ore contains silica, or clay, limestone is used; 
but if the ore contains limestone, some silicate is used. 
The limestone combines with the silicate to form the fus- 
ible slag, which carries with it some of the other impuri- 
ties, and also serves to protect the reduced iron from 
oxidation. When the iron becomes fused, it runs away 
from the slag to the bottom of the hearth, from whence 
it is drawn off from time to time. 



LECTURE NOTES 335 

442. The Chemistry of the Blast Furnace. The chem- 
ical changes which take place in the blast furnace are 
numerous, and some of them are quite complex. They 
differ in different parts of the furnace, and our knowledge 
of some of them, is far from complete. The following are 
some of the more important changes which take place. 

When the hot blast comes in contact with the hot car- 
bon, it forms carbon monoxid, CO. As the ore works its 
way down the furnace, it reaches a point where at a tem- 
perature of 600° to 900°, it meets the hot carbon mon- 
oxid, and is partially reduced, forming a porous mass of 
metalnc iron, thus: 

Fe 2 3 + 3 CO = Fe 2 + 3 C0 2 . 

This action takes place in the upper part of the furnace, 
in what is called the zone of reduction. 

The reduction which takes place here is not quite com- 
plete, but as the ore sinks lower in the furnace it reaches 
the widest part, where the temperature is about 1000°. 
Here the reduction is completed by the hot carbon, and 
the iron begins to take up carbon, forming cast iron. 
This it continues to do until it becomes saturated, sinking 
lower in the furnace all the time; and when it reaches a 
point where the temperature is about 1400°, it melts, and 
runs to the bottom. At this temperature the phosphates 
and some silicates, are reduced, and the phosphorus and 
silicon are taken up by the iron. Some sulfur is taken up 
in the zone of reduction, so that these three elements, to- 
gether with carbon are always found in greater or less 
quantities in cast-iron. The cast-iron is drawn off from 
the hearth and run into sand moulds, where it solidifies, 
forming what are called pigs. 



336 LECTURE NOTES 

443. Cast-iron. Cast-iron is the iron as it comes from 
the blast furnace. It contains from 1.5 to 5 per cent, of 
carbon, and small but varying quantities of sulfur, silicon, 
and phosphorus. When the iron is combined with the 
carbon, it is called white cast-iron; but when the iron con- 
tains the carbon as minute crystals of graphite, it is 
called gray cast-iron. There are several intermediate 
conditions which are known as mottled cast-iron. 

If the iron ore contains manganese, this is reduced with 
the iron, and such iron may contain as much as 6 per cent, 
of combined carbon. This is called specular iron, or 
spiegel. Iron containing more than 20 per cent, of man- 
ganese is called f err o -manganese. 

Cast-iron melts at from 1400° to 1600°, according to 
the amount of carbon which it contains, and when cooled 
suddenly is very hard and brittle. 

444. Wrought Iron, or Malleable Iron. Wrought iron 
is obtained from cast-iron by the process called puddling. 
This is done in a puddling furnace, which has a flat hearth, 
and a low, arched roof. The hearth is lined with iron ore, 
called fettling, and then filled with pigs of cast-iron. The 
flames, pass over the iron, which is soon melted. 

The melted mass is then well stirred and exposed to 
the action of the fettling, and of the air. This oxidizes the 
carbon to carbon monoxid, the escape of which causes the 
mass to swell up, or "boil," so that this process is often 
called pig boiling. As the carbon is removed, the melt- 
ing point of the mass is raised, so that it soon becomes 
pasty. This is then worked up into lumps, called blooms, 
which are placed under a steam hammer, where the slag 
is squeezed out, and the iron welded into a solid mass. 



LECTURE NOTES 337 

Nearly all of the silicon and a large portion of the phos- 
phorus and sulfur are oxidized, and removed with the 
slag. 

Wrought iron contains less than 0.6 per cent, of carbon, 
and usually only traces of the other impurities. It melts 
at from 1900° to 2100° according to its purity, and is 
quite malleable. 

445. Steel. This important variety of iron was form- 
erly denned as iron which contained from 0.6 to 1.5 per 
cent, of carbon, and was capable of being hardened and 
tempered. This definition has now been extended so as 
to include the so-called mild steel, which sometimes con- 
tains less than 0.15 per cent of carbon, and cannot be 
hardened. 

According to our present use of the term, steel includes 
all malleable alloys of iron and carbon which have been 
produced in the molten condition. The difference between 
wrought iron and steel is, therefore, one of structure 
rather than of composition. All steel, having at some- 
time been melted, is much more homogeneous than wrought 
iron. 

The different purposes for which the steel is to be used, 
determines the amount of carbon which it contains. Steel 
rails contain about 0.45 per cent, of carbon ; ordinary 
drills and tools contain from 0.6 to 1 per cent, of carbon; 
while the finest tools contain from 1 to 1.5 per cent, of 
carbon. More than this amount of carbon makes the 
steel very brittle, and it gradually passes into cast iron. 
Hardening steel increases its tenacity, but decreases its 
ductility. 

There are several processes for making steel; it may 



338 LECTURE NOTES 

be obtained from wrought iron by the addition of 
carbon, or from cast iron by the removal of carbon. 
Nearly all steel is now obtained from cast iron. 

446. The Cementation Process. This process possesses 
little but historical interest at the present time, although 
nearly all steel was formerly made in this way. It con- 
sists in heating bars of wrought iron to redness in contact 
with fine charcoal. From seven to ten days are required 
to complete the operation. The surface of the steel then 
presents a blistered appearance, and so is called blister- 
steel. It can be rendered homogeneous by melting in 
crucibles when it is called cast-steel, and is used for the 
finest kinds of cutlery. Just in what way the carbon is 
taken up by the iron is not surely known. It is quite 
probable however, that a portion of the carbon is absorbed 
from the carbon monoxid, which is always formed, and 
a portion, at least, by diffusion of solid carbon into the 
hot iron. 

447. The Bessemer Process. This process consists in 
oxidizing the carbon and silicon in cast-iron, by means of 
a blast of air forced through the molten metal. The oper- 
ation takes place in a large oval-shaped vessel called a 
converter, which is made of wrought iron plates, and lined 
with a paste made from an infusible siliceous rock called 
ganister. It is mounted on trunnions, so that it can be 
tilted, and has an interchangeable bottom, through which 
pass the tuyeres which admit the blast. 

The converter is placed in a horizontal position, and a 
charge of several tons of molten cast-iron poured in. It 
is then raised to a vertical position, and the blast turned 
on. In a few minutes the impurities are burned away, 
the instant the operation is complete being marked by the 



LECTURE NOTES 339 

sudden disappearance of flames from the mouth of the 
converter. Exactly enough molten spiegel (443) is then 
added to supply the right amount of carbon, the blast is 
turned on long enough to mix the materials thoroughly, 
and the steel is then ready for casting. 

448. The Thomas-Gilchrist, or Basic Process. If the 
cast-iron contains more than a very small percentage of 
phosphorus it is not adapted to the manufacture of steel 
by the Bessemer process in the ordinary way, because in 
this process the phosphorus is not removed, but remains 
unchanged in the steel. But if the converter is lined 
with a paste of lime and magnesia, obtained by calcining 
dolomite, the phosphorus combines with this "basic-lin- 
ing," forming a phosphate of calcium and magnesium, 
which passes off as slag. The operation is almost exactly 
like the ordinary Bessemer process, the only difference 
being, that the blast is continued for a short time after 
the flames disappear from the mouth of the converter. 
In this way steel can be made from iron containing a large 
amount of phosphorus. The slag is also valuable, being 
ground, and used as a fertilizer. 

449. The Siemens-Martin Process. This process known 
also as the open hearth process, is, in importance, sec- 
ond only to the Bessemer process. The steel is made 
by melting cast-iron in a reverberatory furnace, and then 
adding the requisite amount of scrap wrought iron, or 
iron ore, or both. The fuel used for heating the furnace 
is gas, which has previously been heated in a Siemens re- 
generative furnace. The furnace hearth is lined with an 
acid, or basic lining, according to the amount of phos- 
phorus in the iron, and the product is much like the 



340 LECTURE NOTES 

Bessemer steel. There are several modifications of the 
open hearth process. 

450. Special Properties of Steel. When a piece of steel 
heated to redness is suddenly cooled, it becomes exceed- 
ingly hard and brittle, and almost perfectly elastic. These 
properties can be almost entirely removed by allowing the 
hot steel to cool slowly, a process called annealing. A 
number of intermediate conditions can be obtained by 
heating the hardened steel to different temperatures, and 
cooling quickly. This process is called tempering. The 
proper temperature is indicated by the color which the 
surface of the steel assumes, and is determined by the 
purpose for which the steel is to be used. The higher 
temperatures produce greater elasticity and brittleness ; 
the lower temperatures, greater hardness and toughness. 

At about 225° steel becomes of a light straw color, and 
when tempered at this point is used for surgical instru- 
ments, and razors. At 240° it is yellow, and used for 
penknives. At 250° it is brownish-yellow and used for 
cold chisels, and metal shears. At 265° shades of purple 
begin to appear, and this temper is used for knives, and 
axes. At 275° it is purple, and used for table knives, 
and scissors. At 285° it is light-blue, and used for swords, 
and watch springs. At 300° and above it is dark-blue, 
and used for chisels, saws, and files. 

When soft iron, or mild steel, is heated for some time 
in contact with powdered charcoal, the outer surface 
becomes highly carbonized. This is known as case-hard- 
ening. 

This process is now applied to armor plates, under the 
name of the Harvey process. The plates are usually made 



LECTURE NOTES 341 

from nickel-steel, and by this treatment the resistance 
which they offer to projectiles is very greatly increased- 

If two pieces of wrought iron or steel are heated to a 
white heat, and then hammered together, they can be made 
to unite and form one solid piece. This process is called 



COMPOUNDS OF IRON 

451. The Ferrous and Ferric Compounds. Iron forms 
two very distinct series of compounds which are known 
as the ferrous and ferric compounds. The ferrous com- 
pounds are quite analogous to the manganous compounds, 
and the iron in them is bivalent, while the ferric com- 
pounds are much like those of aluminum and chromium, 
the iron in them being trivalent. 

Iron dissolves easily in most dilute acids, forming fer- 
rous salts, and hydrogen, the latter having a peculiar 
odor, due to hydrocarbons which are liberated at the same 
time. When dissolved in nitric acid, which is an oxidiz- 
ing agent, unless the acid is very dilute, iron forms a 
ferric compound (See 440). 

The ferric compounds are much more stable than the 
ferrous, and the latter, especially in solution, gradually 
change to the former when exposed to the air, or in the 
presence of other oxidizing agents. On the other hand 
ferric compounds are easily changed to ferrous, in the 
presence of reducing agents. 

THE FERROUS COMPOUNDS 

452. Ferrous Oxid, FeO. This compound is best ob- 
tained by the partial reduction of ferric oxid by heating 
the latter to 300° in a stream of hydrogen. It is a black 



342 LECTURE NOTES 

powder which is easily oxidized when warmed in the air. 

Ferrous hydroxid, Fe(GH) 2 , which corresponds to this 
oxid, is formed when a soluble hydroxid is added to a 
solution of a ferrous salt. It is a white compound, but, 
being very easily oxidized, it is almost never obtained 
white, but has a dirty-green color which gradually changes 
to a red-brown, forming ferric hydroxid. 

453. Ferrous Chlorid, FeCl 2 . This compound is best 
formed by heating iron filings in a stream of hydrochloric 
acid gas. It forms a white crystalline mass which fuses 
at a red heat. 

When iron is dissolved in hydrochloric acid, or the 
anhydrous salt is dissolved in water and the solution con- 
centrated, bluish green crystals, having the composition 
FeCl 2 , 4 H 2 0, are formed. Ferrous chlorid forms double 
salts with the alkaline chlorids. 

454. Ferrous Sulfate, FeS0 4 . This important compound 
was known to the early alchemists. It is also known by 
the commercial names copperas, and green vitriol. It is 
manufactured on a large scale by the oxidation of pyrite 
in the presence of air, and by dissolving iron in sulfuric 
acid. It forms large, green, monoclinic crystals, with 
seven molecules of water, in which respect it resembles 
the corresponding salts of the magnesium group. The 
crystals are somewhat efflorescent, and lose six molecules 
of water at 100°, forming a white powder. With the 
alkaline sulfates it forms monoclinic double salts (328). 
The ammonium double salt, Fe(NH 4 ) 2 (S0 4 ) 2 , 6 H 2 0, — also 
called Mohr's salt — is very stable, and so is much em- 
ployed in volumetric analysis. 



LECTURE NOTES 343 

When ferrous sulfate is heated in the air it decomposes, 
thus : 

2 FeS0 4 = Fe 2 O s + S0 3 + S0 2 . 

When nitric oxid, NO, is passed through a solution of 
ferrous sulfate, a peculiar compound, having the compo- 
sition FeSOJSTO, is formed. This compound is very un- 
stable and decomposes readily on heating (Compare 131). 

Ferrous sulfate is much used in the manufacture of ink, 
and in dyeing. 

455. Other Ferrous Compounds. Some of the other 
ferrous compounds are found in nature although only a 
few of them are of special importance. 

Ferrous sulfid, FeS, is formed by heating together iron 
and sulfur. It is a dark-gray, metallic looking mass, 
which is extensively used in the laboratory for the prepa- 
ration of hydrogen sulfid. It is also formed by treating 
a soluble salt of iron with an alkaline sulfid. 

Ferrous carbonate, FeCO s , occurs in nature as the min- 
eral siderite. It is somewhat soluble in water containing 
carbon dioxid, and so is present in some natural waters. 
When sodium carbonate is added to a solution of a ferrous 
salt, a white precipitate of ferrous carbonate is formed. 
This quickly oxidizes in the air, losing carbon dioxid and 
forming ferric hydroxid. 

THE FERRIC COMPOUNDS 

456. Ferric Oxid, Fe 2 3 . This compound is found in 
nature as the mineral hematite, which is the most valuable 
ore of iron. It can be made artificially by heating ferric 
hydroxid, also by heating ferrous sulfate or oxalate in 
contact with air. The artificial compound is known by 



344 LECTURE NOTES 

many different names, such as rouge, crocus, colcothar, 
and caput mortuum. It is used as a pigment, and for 
polishing glass and metals. 

Ferric hydroxid, Fe(OH) 3 , which corresponds to this 
oxid, is formed when a soluble hydroxid is added to a 
solution of a ferric salt. It is a red-brown compound, 
which, when boiled for some time in water, is converted 
into a number of other hydroxids, all of which are found 
in nature. The more important of these are the miner- 
als xanthosiderite, Fe 2 0(OH) 4 , limonite, Fe 4 3 (OH) 6 , and 
gbthite, FeO,OH. 

457. Ferroso-ferric Oxid, Fe 3 4 . This compound is 
found in nature as the mineral magnetite, which is also a 
valuable iron ore. It can be formed artificially by burn- 
ing iron in oxygen, or by passing steam over red-hot iron. 
It is a black compound which is attracted by a magnet, 
and so is often called the magnetic oxid of iron. The 
mineral magnetite is itself sometimes a magnet, and is 
then called loadstone. 

One of the ferric hydroxids, FeO,OH, like the corre- 
sponding hydroxids of aluminum and chromium, acts like 
a weak acid, forming salts with certain of the metals. 
The magnetic oxid is believed to be a ferrous salt of this 
compound, having the composition (FeO,0) 2 Fe. 

458. The Other Ferric Compounds. Iron forms a large 
number of ferric compounds, most of which are not of 
especial importance. 

Ferric chlorid, FeCl 3 , is formed when iron is dissolved 
in aqua regia, or when chlorin is passed into a solution of 
ferrous chlorid. When the solution is evaporated, a yel- 
low crystalline mass remains, which has the composition 



LECTURE NOTES 345 

2 FeCl s , 3 H 2 0. It is easily soluble in water, alcohol, 
and ether. 

Anhydrous ferric chlorid is formed by heating iron in a 
stream of chlorin. It sublimes quite easily, forming a 
very dark crystalline mass, which, at a high temperature, 
decomposes into ferrous chlorid and chlorin. 

Ferric sulfid, FeS 2 , is found in nature as the mineral 
pyrite, or iron pyrites. It has a golden-yellow color and 
crystallizes in regular forms, the cube being the predom- 
inant form. 

Ferric sulfate, Fe 2 (S0 4 ) 3 , is prepared by dissolving fer- 
ric oxid or hydroxid in sulfuric acid. It forms alums 
with the alkaline sulfates (See 361). 

459. The Cyanogen Compounds of Iron. When potas- 
sium cyanid, KCN, is added to a solution of ferrous, or 
ferric salts, the cyanids, which are at first formed, and 
which are very unstable, dissolve in an excess of the 
potassium cyanid forming double cyanids. These have 
the composition Fe(CN) 2 ,4KCN, and Fe(CN) 8 , 3 KCN, 
respectively. When strong solutions of these double salts 
are treated with acids, hydrogen compounds separate out, 
which are complex acids. These are called hydroferro- 
cyanic, and hydroferricyanic acids, respectively. The 
former has the composition H 4 Fe(CN) 6 , and the latter 
H 3 Fe(CN) 6 . The names "ferro", and " ferri ", are con- 
tractions for ferrous and ferric. Each of the acids 
forms a series of salts, which are called ferrocyauids, and 
ferricyanids, respectively. The iron and cyanogen in these 
compounds cannot be detected by the ordinary reagents, 
and so they are regarded as forming a compound radical, 
in which they act like the halogens. 



346 LECTURE NOTES 

460. Potassium Ferrocyanid, K 4 Fe(CN) 6 . This impor- 
tant compound, which is known commercially as yellow 
prussiate of potash, can be prepared by treating iron, or 
its oxids, with potassium cyanid. It is manufactured on 
a large scale by fusing crude potassium carbonate, and 
adding a mixture of iron filings and nitrogenous animal 
matter, such as horn, hair, blood, and leather. The car- 
bon and nitrogen, together with the potassium carbonate, 
form potassium cyanid, and this acting upon the iron gives 
the ferrocyanid. 

Potassium ferrocyanid forms large yellow crystals, hav- 
ing three molecules of water, and which are easily soluble 
in water. When strongly heated it decomposes thus : 

K 4 Fe (CN) 6 = 4 KCN + N 2 + FeC 2 . 

When heated with concentrated sulfuric acid it decom- 
poses, thus : 

K 4 Fe(CN) 6 + 6 H 2 S0 4 + 6 H 2 = 

FeS0 4 -f 2 K 2 S0 4 + 3 (NH 4 ) 2 S0 4 + 6 CO. 

It is an important reagent in the laboratory, forming 
compounds with a large number of the metals. When 
added to a copper solution it forms copper ferrocyanid 
Cu 2 Fe(CN) 6 , called also Hatchetfs Brown, and added to a 
ferric solution it forms ferric ferrocyanid, Fe 4 [Fe(CN) 6 ] s , 
called also Prussian Blue. 

461. Potassium Ferricyanid, K 3 Fe (CN) 6 . This com- 
pound is not so important as the ferrocyanid. It is 
formed by leading chlorin into a solution of potassium fer- 
rocyanid, thus : 

2 K 4 Fe(CN) 6 + Cl 2 = 2 K 3 Fe(CN) 6 + 2 KC1. 

Potassium ferricya.nid is known commercially as red 



LECTURE NOTES 347 

prussiate of potash. It crystallizes in dark red prisms 
which are soluble in water. With ferrous solutions it forms 
ferrous ferricyanid: Fe s LFe(CN) 6 ] 2 , which is a dark blue 
compound, and known commercially as TurnhulVs Blue. 

Nickel, Symbol, Ni. Atomic Weight, 58.7 

462. History and Occurrence. This metal was discov- 
ered in the mineral niccolite in 1751, by Cronstedt. The 
ore, which was first mentioned in 1694, was called k upfer- 
nic/cel, or false copper, because, while it had the appear- 
ance of a copper ore, no copper could be obtained from it. 

Nickel is always found in meteoric iron. With this ex- 
ception it is never found native. Its most important ores 
are niccolite, NiAs, gersdorffite, NiAsS, and garnierite, 
H 2 (Ni,Mg)Si0 4 . It is also found in connection with cer- 
tain iron and copper ores, usually in the form of millerite, 
NiS, and in several rare minerals. Nickel ore nearly 
always contains some cobalt, and often antimony and 
bismuth. 

463. Preparation. Nearly all the nickel of commerce is 
obtained either from garnierite, or from the sulfids and 
arsenids of nickel which occur in connection with certain 
iron and copper ores. To obtain the nickel from these 
ores they are first roasted, then smelted in a blast furnace 
with lime and slag, and afterwards subjected to a strong 
blast in a silica lined, Bessemer converter. Here most 
of the sulfur, arsenic, and iron, are oxidized, the latter 
being removed in the slag, leaving a matte rich in nickel 
and copper. This is very often reduced with carbon, and 
the nickel-copper alloy used in the manufacture of Ger- 
man silver. To remove the nickel, the matte is fused 
with sodium sulfate and coke, forming sulfids, from which 



348 LECTURE NOTES 

the nickel compound, being heavier and more difficultly 
fusible, is easily separated. The nickel sulfid is then 
roasted and converted into the oxid, and the oxid reduced 
by heating with carbon. 

The matte maybe ground and dissolved in hydrochloric 
acid, from which solution the iron is precipitated with 
ground chalk, and the copper with hydrogen sulfid. The 
solution which now contains onlv the nickel and cobalt, 
is treated with a solution of bleaching powder which pre- 
cipitates the cobalt, after which the nickel is precipitated 
with sodium carbonate, ignited, and reduced with carbon. 

To obtain nickel from garnierite, the ore is smelted in 
a low blast furnace with coke and gypsum, the matte 
roasted and fused with sand and slag, until all the iron 
has been removed, the nickel sulfid converted into oxid by 
roasting, and the oxid reduced with carbon. 

464. Properties and Uses. Nickel is a lustrous, white 
metal which has a grayish tinge when compared with sil- 
ver. It is very hard, takes a high polish, and is the most 
tenacious of all metals. It is quite ductile and malleable, 
although when reduced with charcoal the metal takes up 
some carbon, like cast iron, and this renders it somewhat 
less malleable. It has a specific gravity of 8.9, and melts 
at a somewhat lower temperature than iron. Like iron 
it is attracted by a magnet, and can be welded. 

Nickel is permanent in the air at ordinary temperatures ; 
it decomposes steam at a red heat. It dissolves in the 
common dilute acids, but, if dipped in concentrated nitric 
acid, it becomes "passive," as does iron. 

Nickel is used for making certain alloys and for electro- 
plating iron and steel articles. Argentan, or German 



LECTURE NOTES 349 

silver, is made of two parts of copper, one of zinc, and 
one of nickel. Nickel coins are three parts copper, and 
one part nickel. Nickel added to steel very materially 
increases its toughness, and tensile strength, so that 
nickel steel is an article of great technical importance. It 
is now very largely used for armor plates, and contains 
from three to fifteen per cent, of nickel. 

In electro-nickel plating the solution used is nickel 
ammonium sulfate, the positive electrode being a plate of 
pure nickel, the article to be plated forming the negative 
electrode. 

COMPOUNDS OF NICKEL 

465. The Oxids of Nickel. There are but two well-de- 
fined oxids of nickel. These are : 

Nickel monoxid, NiO, and 
Nickel sesquioxid, Ni 2 O s . 

Other oxids have been described, but their existence is 
doubtful, and they are usually regarded as mixtures of 
the oxids, or of nickel and the oxids. 

Nickel monoxid,W\0, is found in nature as the mineral 
bunsenite, and may be formed by heating nickel carbonate, 
or hydroxid, out of contact with the air. It is a green 
crystalline substance, which dissolves easily in acids form- 
ing nickel salts, and is easily reduced to the metal with 
hydrogen, carbon monoxid, or carbon. 

Nickel sesquioxid, ^N\fi 3 is formed; by carefully igniting 
nickel nitrate, or carbonate, in the air. It is a black 
powder, which^at a higherjtemperature is converted into 
the monoxid. With acids it acts like a peroxid (232). 

Nickel hydroxid, Ni(OH) 2 , is formed by adding solu- 



350 LECTURE NOTES 

ble hydroxid to a solution of a nickel salt. It is a pale 
green substance, which is easily soluble in acids, forming 
nickel salts. It is also soluble in ammonium hydroxid. 

466. Other Compounds of Nickel. Nickel forms only 
one series of compounds. These correspond to the mon- 
oxid, the nickel in them being bivalent. 

Nickel chlorid, NiCl 8 , is formed by dissolving the oxid 
or carbonate in hydrochloric acid. It forms green, pris- 
matic crystals, with six molecules of water. 

Nickel sulfide NiS, is found in nature as the mineral 
millerite, in the form of brass -yellow, hair-like or needle 
shaped crystals. It is formed as a black amorphous pow- 
der, by the addition of an alkaline sulfid to a solution of a 
nickel salt. 

Nickel sidfate, NiS0 4 , is formed by dissolving the metal, 
its oxids, or carbonate, in sulfuric acid. It forms green 
crystals with seven molecules of water, which are isomor- 
phous with magnesium sulfate, and which form monoclinic 
double salts (328). The nickel-ammonium sulfate is most 
important, being employed in the process of nickel-plating. 

Nickel nitrate, Ni(NO ) 2 , forms green tabular crystals 
with six molecules of water of crystallization. 

Nickel carbonyl, Ni(CO) 4 , is an interesting compound, 
formed by passing carbon monoxid over finely divided 
nickel heated to about 50°. It is a liquid, boiling at 43°, 
and solidifying at — 25°. When the vapor is heated to 
60° it explodes, but if mixed with some indifferent gas, 
such as hydrogen or nitrogen, and passed through a tube 
heated to 180°, it is decomposed, pure nickel being depos-. 
ited, and carbon monoxid liberated. 



LECTURE NOTES 351 

Cobalt. Symbol, Co. Atomic Weight, 59. 

467. History and Occurrence. This metal was first 
obtained in 1735, by Brandt. Specimens of Egyptian 
glass, as well as G-reek and Roman pigments which con- 
tained cobalt, have been found, and ores, now known to 
contain cobalt, have been used for the purpose of coloring 
glass, since the middle of the 16th century. 

Cobalt has been found in small quantities in meteorites, 
with which exception it does not occur free in nature It 
is almost always found in small quantities in nickel 
ores. The most important cobalt minerals are linnaeite, 
(Co, Ni) 3 S 4 , cobaltite, CoAsS, and smaltite, (Co, Ni) As 2 . 

468. Preparation, Properties, and Uses. The cobalt 
ores are treated in the same way as the ores of nickel. 
After separating the other metals as completely as possi- 
ble, the cobalt-nickel matte is dissolved in hydrochloric 
acid, any impurities which may remain are precipitated, 
and the cobalt separated from the nickel by precipitating 
it as cobaltic hydroxid, Co(OH) 3 , by the careful addition 
of bleaching powder. The cobaltic hydroxid is changed 
to the oxid, and the oxid reduced to the metal, by ignit- 
ing in a stream of hydrogen, or with carbon. 

Cobalt is a lustrous, white metal, with a distinct red- 
dish shade. It is harder than iron, is quite malleable, 
and very tenacious. It is magnetic, but less so than iron. 
Its specific gravity is 8.5 to 8.9, and it melts at a some- 
what lower temperature than iron. 

Cobalt dissolves in hydrochloric, sulfuric, and nitric 
acids. At the ordinary temperature it is unaltered in the 
air, or water, but at a red heat it decomposes steam, form- 
ing cobaltous oxid, CoO. It does not become passive 
when dipped in nitric acid, as do iron and nickel. 



352 LECTURE NOTES 

The principal compounds of cobalt are those in which 
the metal is bivalent. These are generally stable, and 
isomorphous with the corresponding ferrous salts. The 
hydrate d compounds have a reddish color, while the an- 
hydrous compounds are blue. There are a few cobaltic 
compounds in which the metal is trivalent. 

COMPOUNDS OF COBALT 

469. The Oxids of Cobalt. Cobalt forms three well 
denned compounds with oxygen. These are : 
Cobalt monoxid, or Cobaltous oxid, CoO, 
Cobalt sesquioxid, or Cobaltic oxid, Co 2 3 , and 
Cobalto-cobaltic oxid, Co 3 4 . 

In addition to these, there are several mixtures of the 
first two, having a fairly definite composition. 

Cobalt monoxid, CoO, can be obtained as a light- brown 
powder, by heating the carbonate in the absence of air, 
or by heating the higher oxids to redness in a stream of 
carbon dioxid. 

Cobalt sesquioxid, Co 2 O s , is a dark-brown powder ob- 
tained by gently igniting the nitrate until the red fumes 
cease to be evolved. It dissolves in cold acids forming 
brown solutions which contain unstable cobaltic salts. 
When warmed with acids, it acts like a peroxid. 

Cobalto-cobaltic oxid, Co 3 4 , is formed as a black pow- 
der, when either of the other oxids is strongly heated in 
the air. 

Two hydroxids are also known, which correspond to 
the first two oxids, and which may be obtained by pre- 
cipitating the corresponding salt solutions with a soluble 
hydroxid. When sodium hydroxid is added to a cold 



LECTURE NOTES 353 

solution of a cobaltous salt, a blue basic salt is form- 
ed, which on boiling is changed to the red cobaltous hy- 
droxid, Co(OH) 2 . On exposure to the air, this gradually 
oxidizes, and becomes brown, by the formation of cobal- 
tic hydroxid, Co(OH) 3 . 

470. The Cobaltous Compounds. A large number of 
these compounds exist, but only a few of them are impor- 
tant. 

Cobaltous chlorid, or cobalt chlorid. CoCl 2 , is formed 
by dissolving the metal, its carbonate, or oxids, in hydro- 
chloric acid. It forms dark-red, prismatic crystals, with 
six molecules of water. These, upon heating, gradually 
lose water, and at 120° become anhydrous, and blue in 
color. On exposure to the air, the anhydrous compound 
takes on water, and becomes red. It is to this property 
that cobaltous chlorid owes its use in the so-called sym- 
pathetic inks. 

Cobaltous sulfate, or cobalt sulfate, CoS0 4 , is obtained 
by the usual methods, and forms dark-red crystals with 
seven molecules of water. It is isomorphous with ferrous 
sulfate, and forms monoclinic double salts with the alka- 
line sulfates. 

Cobalt nitrate, Co(N0 3 ) 2 , forms a red, deliquescent, crys- 
talline mass, which is much used in the laboratory, espe- 
cially in blowpipe analysis. 

Cobalt silicate, or smalt, is a beautiful blue pigment, 
which is made commercially by fusing cobalt ores with 
sand and potassium carbonate. The glass thus formed is 
powered fine and used as a pigment, for which purpose it 
is highly valued, being unchanged by light, and unaffected 
by acids and alkalies. 



354 LECTURE NOTES 

471. The Cobaltic Compounds. When cobaltic oxid, 
or hydroxid is dissolved in cold, dilute acids, a brown 
solution is formed which contains a cobaltic salt corre- 
sponding to the acid used. These salts are all unstable, 
and are more or less completely decomposed by heating, 
with formation of the corresponding cobaltous salt. They 
also form, in some cases, double salts with the correspond- 
ing alkaline salts, which are quite stable. Thus, when 
potassium nitrite, KNO z , is added to an acetic acid solu- 
tion of a cobaltous salt, a yellow, crystalline precipitate 
is slowly formed, which is cobaltic potassium nitrite, 
Co(N0 2 ) 3 , 3 KN0 2 . Cobaltic sulfate forms a series of un- 
stable alums with the alkaline sulfates. 

Cobalt forms cobaltocyanids and cobalticyanids which 
are analogous to the ferro- and ferricyanids, and which 
may be regarded as salts of a complex acid containing 
cobalt. 

472. The Cobaltamines. If ammonium hydroxid is 
added to a solution of cobalt chlorid, the precipitate first 
formed dissolves in excess of the ammonium hydroxid, 
and if this solution is left standing for some time exposed 
to the air, it becomes first brown, and finally red. If hy- 
drochloric acid is now added to the solution, a brick-red, 
crystalline powder is precipitated. This is called roseo- 
cobaltic chlorid, and has the composition CoCl 3 ,5 NH 3 ,H 2 0. 
If the red solution is boiled before precipitating, a red 
powder, called pupureo -cobaltic chlorid, CoCl 3 , 5 NH 3 , is 
formed. If ammonium chlorid is present in the solution 
a brown compound called I uteo -cobaltic chlorid, CoCl 3 , 6NH 3 , 
is precipitated. The other cobalt salts treated in a 
similar way yield similar compounds. These are known 
as the ammonia- cobalt compounds, or the cobaltamines. 



LECTURE NOTES 355 

473. The Platinum Metals. The metals which consti-' 
tute the second and third sub-divisions of connecting 
elements in division B of Group VIII, are very closely 
related to each other. They all occur free in nature, 
associated together in what is called platinum ore, and so 
are commonly called the platinum metals. The second sub- 
division consists of the metals ruthenium, rhodium, and 
palladium, and is often called the ruthenium group, or the 
light platinum metals, the third sub-division, which con- 
sists of the metals osmium, iridium, and platinum, being 
called the platinum group, or the heavy platinum metals. 

Platinum ore is found in small grains, or nuggets, in 
alluvial deposits, and in river sand, the principal localities 
being Brazil, Australia, Borneo, the Ural Mts., and Cali- 
fornia. It usually contains from 60 to 85 per cent, of 
platinum, the remainder consisting of the other platinum 
metals, together with varying quantities of gold, copper, 
and iron. The minerals called platiniridium, an alloy of 
platinum and iridium, and osmiridium, an alloy of osmium 
and iridium, are both found with the platinum ore, and 
contain the other metals of the group in small quantities. 

The platinum metals all fuse with great difficulty. 
Ruthenium and osmium are quite easily oxidized when 
heated in the air, the others remaining unchanged. Pal- 
ladium dissolves in nitric acid. Aqua regia dissolves 
platinum and also the other metals when alloyed with 
much platinum, but when the metals are pure it is with- 
out action upon them. They each form several oxids, and 
series of salts corresponding to these. Ruthenium and 
osmium are not so basic as the others, and exhibit a non- 
metallic character in forming salts corresponding to hypo- 
thetical acid-forming oxids. They are also remarkable in 



356 LECTURE NOTES 

each forming a volatile tetroxid. Each of the platinum 
metals forms complex compounds with ammonia, and with 
cyanogen. 

474. The Metallurgy of the Platinum Metals. The 

process of separating the platinum metals is quite com- 
plicated, and depends mainly upon the solubility of the 
different metals in aqua regia. The ore, after separating 
it as much as possible from earth and sand, is treated 
with dilute aqua regia. This dissolves gold, copper, and 
iron. The remainder is then treated with concentrated 
aqua regia which dissolves the platinum and palladium, 
together with some iridium, and traces of ruthenium and 
rhodium. The residue consists of the undissolved portion 
of these latter metals together with an alloy of osmium 
and iridium. 

Ammonium chlorid is then added to the solution and 
platinum and iridium precipitated as double chlorids. The 
amount of iridium is usually so small that it is not sepa- 
rated from the platinum. Upon igniting the double salts, 
the platinum and iridium are left as a spongy mass, which 
is either ignited, and hammered out, or fused in the oxy- 
hydrogen flame, and run into moulds. 

Ruthenium. Symbol, Ru. Atomic Weight, 101.7. 

475. Ruthenium and its Compounds. This element was 
discovered in 1845 by Claus. It occurs in platinum ore, 
in osmiridium, and as the mineral laurite, Ru 2 S 3 . It is 
usually obtained from the "platinum residue," which is 
the undissolved portion remaining after treating the plat- 
inum ore with concentrated aqua regia. The process of 
extracting the ruthenium from this residue is long and 
complicated, and will be omitted. 



LECTURE NOTES 357 

Ruthenium is a greyish-white metal, having a specific 
gravity of 12.27, and is, next to osmium, the most diffi- 
cultly fusible of all known metals. It is insoluble in all 
acids, and only very slowly attacked by aqua regia. When 
fused with potassium hydroxid and an oxidizing agent, it 
forms potassium ruthenate, K 2 Ru0 4 , which is soluble in 
water. Perhaps the most interesting of ruthenium com- 
pounds is the tetroxid, Ru0 4 , which is formed by heating 
the metal to 1000° in a stream of oxygen, or by passing 
chlorin into a solution of potassium ruthenate. It is a 
yellow crystalline mass which melts at 25.5°, and can be 
distilled under slightly reduced pressure. 

In addition to the complex ammonia, and cyanogen com- 
pounds, common to the metals of Group VIII, ruthenium 
forms a nitrosochlorid, RuCl 3 NO, which forms a series of 
salts with the alkaline chlorids. The potassium com- 
pound has the constitution RuCl 3 NO, 2 KC1, written also 
K 2 RuCl 5 NO, and called potassium nitroso-chlororuthenate. 

Rhodium. Symbol, Rh. Atomic Weight, 103. 

476. Rhodium and its Compounds. This metal was 
discovered in 1 803 by Wollaston in the ' 'platinum residue. " 
It is a bluish-white metal with a bright luster and re- 
sembles aluminum. It has a specific gravity of 12.1 and 
is more easily fusible than ruthenium. 

Rhodium forms several oxids, all of which are basic. 
Rhodium chlorid, RhCl 3 , forms rose-colored double salts 
with the alkaline chlorids. Its ammonia compounds most 
resemble those of cobalt. 

Palladium. Symbol, Pd. Atomic Weight, 107. 

477. Palladium and its Compounds. This metal was 



358 LECTURE NOTES 

discovered in some platinum ore from Brazil, in 1803, by 
Wollaston. It occurs in nature in most platinum ore, 
also in gold ore from the Hartz Mts., and from several 
places in South America. It is obtained from the palla- 
dium-gold alloy by fusing with silver, and treating this 
with nitric acid, when the silver and palladium dissolve, 
leaving the gold. The silver is then precipitated with 
sodium chlorid, and the palladium with zinc as the metal, 
or as cyanid with potassium cyanid,. which by ignition 
leaves the metal. 

Palladium is a silver-white metal, which has a specific 
gravity of 11.4, and melts at about 1400°. When heated 
to redness it assumes a violet color, but at a higher tem- 
perature becomes white again. It forms two series of 
compounds, palladious and palladic. In the former the 
metal is bivalent, and in the latter it is quadrivalent. 

478. Palladium and Hydrogen. One of the most interest- 
ing properties of palladium is its " occlusion" of hydrogen, 
discovered by Graham in 1866. If in the decomposition 
of water by the electric current, a piece of palladium is 
used for the negative electrode, the hydrogen, which is 
set free from the water, is absorbed by the palladium, so 
that until the metal is saturated, no hydrogen is given 
off. In this way palladium can be made to absorb 935 
volumes of hydrogen (Compare 45). The amount absorb- 
ed, however, depends upon the temperature: at all ordi- 
nary temperatures, some absorption takes place, but this is 
greatest at about 100°. 

For many reasons Graham believed that the hydrogen 
formed an alloy with the palladium. It is believed by 
some, that this alloy is a definite compound, Pd 2 H, but 
this has not been confirmed. The palladium increases 



LECTURE NOTES 359 

about 10 per cent, in volume, the alloy having a specific 
gravity of 11.06. Sometimes, in contact with the air, it 
becomes heated, the hydrogen being oxidized to water. 
It begins to decompose at 130°, but is completely decom- 
posed only after long continued heating to redness. The 
hydrogen given off is apparently in the nascent state, 
since it acts as a reducing agent. 

Osmium. Symbol, Os. Atomic Weight, 191. 

479. Osmium and its Compounds. This metal was dis- 
covered in the "platinum residue " in 1804, by Tennant. 
Its most important ore is osmiridium. From these ores 
it may be prepared by strongly heating in oxygen, when 
osmium tetroxid, Os0 4 , is formed. This may be reduced 
to the metal, by passing its vapor, mixed with carbon 
monoxid and carbon dioxid, through a red-hot porcelain 
tube. 

Osmium resembles ruthenium more than any other met- 
al. It has a bluish-white color and is the most difficultly 
fusible of all the metals, being fusible only in the electric 
furnace. Its specific gravity is 22.48, and it is therefore 
the heaviest of all metals. Its most interesting and im- 
portant compound is the tetroxid, Os0 4 , which is com- 
monly, though improperly, called osmic acid. This com- 
pound boils without decomposition at a little above 100°, 
and is used in microscopy for hardening and coloring the 
tissues. Some of the compounds of osmium are exceed- 
ingly poisonous. 

Iridium. Symbol, Ir. Atomic Weight, 1 93.1. 

480. Iridium and its Compounds. This metal was 
discovered in the "platinum residue" in 1804, by Ten- 



360 LECTURE NOTES 

nant. It also occurs in platiniridium and osimiridium. In 
the separation of the platinum metals, it is found with 
the platinum, or with the osmium. From the latter the 
osmium is separated as given under osmium (479), leaving 
the iridium, while from the former it is separated by 
forming the trichlorid, IrCl 3 , which forms an easily solu- 
ble salt with ammonium chlorid. 

It is a grayish-white metal, which is much harder than 
the other platinum metals. It is nearly as heavy as os- 
mium, its specific gravity being 22.42. It is, next to 
osmium and ruthenium, the most difficultly fusible of all 
the metals. It is used, both alone and alloyed with 
osmium, for pointing gold pens. Alloyed with 9 parts of 
platinum, it is used for making standard meter bars, and 
for making electrodes to be used in corrosive liquids. If 
platinum is alloyed with a small amount of iridium, it is 
very much improved for many purposes. 

Platinum. Symbol, Pt. Atomic Weight, 194.9. 

481. History, Occurrence, and Preparation. This very 
important metal was probably first observed in the six- 
teenth century ; for a writer of that period states that a 
metallic substance had been found in a Mexican mine, 
which could not be fused. Since platinum is found in 
this same locality, it is probably the metal referred to. 
About two centuries later it was described by de Ulloa. 
In 1823 it was discovered in the Ural Mts. 

It almost always occurs free in nature in the form of 
platinum ore, which has already been described (473). It 
is also found in Canada in the mineral sperrylite PtAs 2 . 
The general process of separating the platinum metals 
has also been described (474). 



LECTURE NOTES 361 

482. Properties. Platinum is a greyish -white metal 
which is very malleable and ductile. It has a specific 
gravity of 21.5, and at a white heat it can be welded. It 
can be fused in the oxyhydrogen flame, its melting point 
being estimated at about 2000°. It can be volatilized in 
the electric arc. 

It is not affected by any single acid, but dissolves quite 
easily in aqua regia. When platinum is precipitated from 
a solution of its salts, it appears as a very fine black 
powder known as platinum black. When certain salts 
such as ammonium chloroplatinate, (NH 4 ) 2 PtCl 6 , or pla- 
tinic chlorid, PtCl , are strongly ignited, platinum re- 
mains in a finely divided form known as platinum sponge. 
This is a little more coherent than platinum black, but 
both forms possess the power of absorbing large quanti- 
ties of hydrogen and oxygen. It acts thus as an oxidiz- 
ing agent, and as such is sometimes used in organic chem- 
istry. 

483. Uses. Platinum is a metal of great use to the 
chemist, both on account of its infusibility, and its resis- 
tance to the action of most chemicals. It is used in the 
form of dishes, crucibles, foil, wire, and tubes, and these 
forms of apparatus make it possible for the chemist to 
carry on successfully much work that was impossible be- 
fore the discovery of this metal. Platinum is, however, 
not indestructible, and care should be exercised in the use 
of it. 

Platinum easily alloys with metals having a low melt- 
ing point, and is more or less attacked when alkaline hy- 
droxids or cyanids, or potassium nitrate, are fused in 
contact with it. Phosphorus and arsenic combine with 



362 LECTURE NOTES 

it when heated, and if a platinum vessel is heated in a 
smoky flame, it becomes brittle from absorption of car- 
bon and the formation of a carbid of the metal. It is best 
cleaned by scouring with wet sea-sand, or by fusing with 
it acid potassium sulfate, or an alkaline carbonate. 

484. The Compounds of Platinum. Platinum forms 
both platinous and platinic compounds. In the former it 
is bivalent, and in the latter quadrivalent. Only a few 
of its compounds are of practical importance. When 
platinum dissolves in aqua regia it forms what is called 
chloroplatinic acid, H 2 PtCl g , which on evaporation of the 
solution, forms crystals with six molecules of water. The 
hydrogen in this compound is replaceable by the metals, 
forming a series of compounds called chloroplatinates, or 
platinochlorids. 

Potassium, chloroplatinate, K 2 PtCl 6 , is formed when a 
potassium salt is added to a solution of the acid. It forms 
reddish-yellow, octahedral crystals, which dissolve in 
water with difficulty, and are insoluble in alcohol. The 
corresponding sodium salt is easily soluble in alcohol, so 
that sodium and potassium are separated by this method, 
in analytical chemistry. 

Ammoriium chloroplatinate, (NH^PtClg, is formed in a 
similar way, and closely resembles the potassium salt. 
On being ignited it decomposes leaving platinum sponge. 

Platinum forms a large number of complex ammonia, 
and cyanogen compounds. 



INDEX 



Absolute temperature, 9 

Acetylene, 151 

Acid, arsenic 135 

arsenious, 134 
boric, 175 
carbonic, 161 
chloric, 65 
chlorous 64 
chlor-auric, 238 
chloroplatinic, 362 
chromic, 317 
dichromic, 317 
hydriodic, 72 
hydrobromic, 68 
hydrochloric, 61 
hydrocyanic, 164 
hydroferricyanic, 345 
hydroferrocyanic, 345 
hydrofluoric, 75 
hydrofluosilicic, 168 
hydrosulfuric, 82 
hypochlorous, 64 
iodic, 73 
metaboric, 176 
metantimonic, 307 
metaphosphoric,126 
metarsenic, 135 
metasilicic, 172 
metastannic, 290 
muriatic, 62 
nitric, 113 

nitro-hydrochloric, 115 
nitrosulfonic, 91 
Nordhausen sulfuric, 93 
ortho-antimonic, 307 
ortho-arsenic, 135 
ortho-phosphoric, 125 



Acid, osmic, 359 
" perchloric, 66 
" periodic, 73 
" phosphoric, 125 
" pyro-antimonic, 307 
" pyro-arsenic, 135 
" pyroboric, 176 
" pyrophosphoric, 126 
" pyrosulfuric, 93 
" radical, 33 
" reaction, 33 
" salts. 35 
" silicic, 171 
" stannic, 290 
" sulfuric, 88 
" sulfurous, 87 
" tetraboric, 176 
" thiosulfuric, 93 

Acids, 32 
" basicity of, 33 

Adhesion, 12 

Affinity, chemical, 22 

Agate, 169 
| Alabaster, 245 
I Alkalies, 34, 187,196 
j Alkali metals, 196 

Alkaline earths, 239 
" reaction, 34 

Allotropic forms, 47 

Alloys, 184 
" important, 185 

Alumina, 275 

Aluminates, 275 

Aluminum, 271 
" bronze, 186 
chlorid, 276 
" group, 270 



364 



INDEX 



Aluminum hydroxid, 275 
oxid, 274 
" silicates, 278 
sulfate, 276 
Alums, 277 

Amalgamation process, 230 
Amalgams, 186, 264 
Amethyst, 169 
Amido group, 108 
Ammonia, 105 
Ammonia-soda process, 205 
Ammonium, 108, 215 
" amalgam, 216 
" carbonate, 219 

chlorid, 217 
" chloroplatinate, 362 

hydroxid, 107, 216 
" molybdate, 319 
" nitrate, 218 
" phospho-molybdate, 319 
" sodium hydrogen phos- 
phate, 219 
sulfate, 218 
sulfids,217 
Analysis, 37 
Anglesite, 291, 296 
Anhydrids, 33 
Anhydrite, 240, 245 
Animal charcoal, 147 
Annealing, 340 
Anode, 223 
Antimonates, 307 
Antimony, 302 

butter of, 307 
halogen comp's. of, 307 
other comp's. of, 309 
oxids of, 305 
oxy-acids of, 306 
trisulfid, 308 
Antifriction metals, 186 
Antichlor, 60 
Apatite, 240, 246 
Aquamarine, 252 



Aqua regia, 115 

Aqueous vapor, tension of, 53 

Argentite, 228 

Argon, 101, 194, 329 

" and its related gases, 101 
Aragonite, 240 
Argentan, 348 
Argol,208 
Argyrodite, 285 
Arsenic, 128 

disulfid, 135 

" halogen comp's. of, 131 

" Marsh's test for, 137 

" mirror, 138 

" oxy-acids of, 134 

" pentasulfid, 136 

" pentoxid, 134 

" Reinsch's test for, 138 
trisulfid, 136 
Arsenious oxid, 132 
Arsenolite, 132 
Arsin, 130 
Atmosphere, 98 

" composition of, 99 

" pressure of, 10 
Atomic attraction, 22 

" heat, 182 

" theory, 4, 24 

" weight, 25 

weights, table of, 21 
Atoms, 4, 19 
Autotype process, (photography) 

318 
Avogadro's law, 15 
Azurite, 221, 227 

Babbit's metal, 186 
Barite, 248 
Barium, 248 
" carbonate, 250 

chlorid, 249 
" nitrate, 250 
" oxids and hydroxid of, 249 
sulfate, 250 



INDEX 



365 



Barometer, 9 
Baryta water, 249 
Base, 34 

Bases, acidity of, 34 
Basic process (steel), 339 

" reaction, 34 
Bauxite, 271 

Belgian process (zinc), -258 
Bell metal, 186 
Beryl, 251 
Beryllium, 251 
Bessemer process (steel), 338 
Binary compounds, 39 
Bismite, 310 
Bismuth, 309 

" nitrate, 312 

" other comp's of, 312 

" - oxids of, 311 
Bismuthinite, 310,^312 
Bittern, 67 

Black ash process (soda), 205 
Black lead, .144 
Blast furnace, 333 

" " chemistry of, 335 

Bleaching, 60 

" powder, 64, 244 
Blister copper, 222 

steel, 338 
Boiling point, :8, 53 
Boort, or bort, 142 
Bone ash, 117 

black, 117 
Borates, 175 
Borax, 176, 207 
Bornite, 221 
Boron, 173 

" compounds of, 174 
Boyle's law, 10 
Brass, 186 
Bra unite, 322 
Britannia metal, 186 
Bromid-gelatin process, (photog- 
raphy), 234 



Bromin, 67 

" oxids and oxy-acids of, 69 
Bronze, 186 
Brookite, 281 
Brucite, 254 
Bunsen lamp, 156 

Cadmia, 261 
Cadmium, 261 

" compounds of, 262 
Caesium, ^215 
Calamine, |257 
Calcite, 162, 240 
Calcium, 239 

carbid, 246 
carbonate, 247 
chlorid, 243 
fluorid, 243 
hydroxid, 241 
oxid, 241 
phosphate, 246 
" sulfate,'; 245 
Caliche, 70 
Calomel, 265 
Calorie, 17 
Calorimeter, 18 
Carat, 142 
Carbon, 139 
" allotropic forms of, 140 
" amorphous, 144 
" dioxid, 158 
" " solid and liquid, 160 

disulfid, 162 
" group, 139 
" halogen comp's. of, 156 
" hydrogen comp's. of, 148 
" monoxid, 157 
Carbonado, 142 
Carbonates,;158, 162 
Carnallite, 208, 252 
Carrie's ice machine, 108 
Carthusian powder, 308 
Case hardening, 340 
Cassiterite, 286, 289 



366 



INDEX 



Cast iron, 336 

Castner's process (sodium), 199 

" " (sodium hydrox- 

id), 200 
Caput mortuum, 344 
Carnelian, 169 
Catalysis, 45 
Cathode, 223 
Caustic potash, 210 

soda, 199 
Celestite, 247 

Cementation process (steel), 338 
Cements, mortars and, 242 
Cerite, 283 
Cerium, 283 
Cerussite, 291 
Chalcedony, 169 
Chalcocite, 221 
Chalcopyrite, 221 
Chalk, 162, 240 
Chalky air, 158, 
Chamber acid crystals, 91 
Charcoal, 146 

" animal, 147 
Chemical action, modes of, 36 
affinity, 22 
changes, 3 

combination, laws of, 23 
equations, 37 
formulas, 31 
nomenclature, 39 
Chemism, 22 
Chili saltpeter, 113, 203 
Chlorargyrite, 228, 232 
Chlorates, 65, 210 
Chlorid of lime, 64, 244 
Chlorin, 58 

" oxids of, 62 

" oxy-acids of, 63 

" " " constitution of 66 

" peroxid, 63 
Chloroform, 156 
Choke damp, 158 
Chromates, 315 



Chrome alum, 278, 318 

" green, 316 

" iron ore, 317 
yellow, 297 
Chromic oxid, 315 
Chromite, 313 
Chromites, 316 
Chromium, 313 

" oxids of, 314 
trioxid, 316 

" sulfate, 318 
Chrysoberyl, 251, 276 
Cinnabar, 263 
Classification of the elements, 28 

" periodic, 29 
Clay, 278 
Coal, 145 

" gas, 152 
Coarse metal (copper), 222 
Cobalt, 351 

•' amines, 354 

" oxids of, 352 
Cobaltic compounds, 354 
Cobaltite, 351 
Cobaltous compounds, 353 
Co-effl cient of expansion(gases),6 
Cohesion, 12 
Coke, 147 
Colcothar, 344 
Colloids, 172 
Columbite, 301 
Columbium, 300 
Combustible, 154 
Combustion, 153 

" supporters of, 154 
Compound radicals, 30 
Compounds, 19 
Condy's 6 liquid, 327 
Constitutional formulas, 31 
Contact action, 45 
Copper, 220 

" compounds of, 225 

" group, 220 



INDEX 



36' 



Copper, other salts of, 227 

oxid, 226 
" refining of, 223 

sulfate, 226 
Copperas, 342 
Coral, 162, 240 
Corosive sublimate, 267 
Corundum, 27], 274 
Crith, 43 
Critical pressure, 102 

" temperature, 102 
Crocoite, 313 
Crocus, 344 
Crookesite, 280 
Cryolite, 74, 271 
Crystallization, water of, 54 
Crystalloids, 172 
Cupric compounds, 225 
Cuprite, 221 

Cuprous compounds, 225 
Cyanid process (gold), 237 
Cyanids, 58, 161 
Cyanogen, 58 

gas, 163 

Dalton's atomic theory, 24 
Davy's safety lamp, 155 
Daguerreotype process, 233 
Deliquescence, 55 
Desiccators 55 
Denitrating tower, 90 
Density, 5 

Departments of science, 2 
Dialysis, 171 
Diamond, 141 

" preparation of, 143 
Diaspore, 271, 275 
Dichromates, 317 
Didymium, 301 
Dolomite, 240, 252 
Double cyanids, 164 
Drummond light, 241 
Dutchjiquid, 150 

" process, (white lead), 297 



Dyads, r -26 

Earths, 238 

rare, 269 
Effect of heat on matter, 6 

" of pressure on matter, 10 
Efflorescence, 54 
Electrotyping, 224 
Electro negative atoms, 29 

" positive atoms, 29 
Elements, 19 

" classification of, 28 

" connecting, 329 

" natural groups of, 190 

" names and symbols of, 19 

" table of, 21 
Emerald, 251 
Emery, 271, 275 
Empirical formulas, 31 
Epsom salts, 255 
Equations, chemical, 37 
Erbium,£70 
Ethylene, 150 
Ferric compounds, 343, 344 

" oxid,; 343 
Ferricyanids, 345 
Ferro-cyanids, 345 
Ferro-manganese, 336 
Ferroso-ferricoxid, 344 
Ferrous chlorid, 342 

" compounds, other, 343 
oxid, 341 
sulfate, 342 
Fettling, 336 
Fire damp, 149 
Fixed air, 158 

" alkali, 196 
Flame, nature of, 155 

" luminosity of, 155 

" oxidizing, 156 

" reducing, 156 
Flint, 169 
Fluorescence, 243 
Fluorin, 74 



368 



INDEX 



Fluorite, 74, 240, 243 

Force, 1 

Formulas, constitutional, 31 

" empirical, 31 
Franklinite, 257 
Freezing point, 8 
French method (white lead), 

Gadolinium, 270 
Gahnite, 276 
Galenite (galena), 291 

" argentiferous. 228 
Gallium, 279 
Galvanized iron, 259 
Garnierite, 347 
Gas carbon, 147 
Gases, diffusion of, 14 

" liquefaction of, 102 
Gay-Lussac tower, 91 
German silver, 348 
Germanium, 285 
Gersdorffite, 347 
Gibbsite, 275 
Glass, 170 
Glauber's.salt, 202 
Glover tower, 90 
Glucinum, 251 
Gold, 235 

" coinage, 237 

" mining, 236 

" compounds of, 238 
Goslarite, 260 
Goethite, 344 
Gram, 5 
Granite* 272 
Graphite, 144 
Gravitation, 4 
Gray cast-iron, 336 
Greenockite, 261 
Green vitriol, 342 
Guignet's green, 316 
Gun metal, 186 
Gunpowder, 213 
Gypsum, 240 



Hall method (aluminum), 272 
Halogen acids, 32 

" group, 57 
Hardness (water), 55 
Harvey process (steel), 340 
Hatchett's brown, 346 
Hausmannite, 322 
Heat, atomic, 182 

" effect of on matter, 6 

" latent, 16 

" measurements 

" specific, 181 

" unit of, 17 
Helium, 102, 193, 329 
Hematite, 332, 343 
Horn silver, 232 
Hydraulic cement, 242 

" lime, 242 
Hydrocarbons, 148 
Hydrogen, 42 

" arsenid, 130 

" bromid, 68 

" chlorid, 61 

" dioxid, 56 

" fluorid, 75 
iodid, 72 

" monoxid, 49 

" occlusion of, 44 

" phosphids, 121 
sulfid, 81 
Hydrogenium, 44 
Hydroxy], 31 
Hyposulfites, 94 

Ignition, temperature of, ]54 
Ilmenite, 281 
Imido group, 108 
Indium, 279 
Iodin, 70 

" oxids and oxy-acids of, 73 
Iridium, 359 
Iron, 331 

" cast, 336 

" compounds of, 341 



INDEX 



369 



Iron cyanogen comp's, of, 345 
magnetic oxid of, 344 
malleable, 336 
metallurgy of, 333 
passive, 333 
pyrites, 345 
wrought, 336 

Isomorphous compounds, 256 

Jasper, 169 
Kainite, 252 
Kaolin, 272 
Kelp, 70, 204 
Kermes, 308 
Kieselguhr, 169 
Kieserite, 252 
King's yellow, 136 
Kohl, 308 

Krypton, 102, 194, 329 
Kupfernickel, 347 
Lakes, 275 
Lampblack, 147 
Lanthanum, 301 
Lapis-lazuli, 228 
Latent heat, 16 

" of water and steam, 52 
Laughing gas, 110 
Laurite, 356 
Law of Boyle, 10 

" " Charles, 6 

" " Dalton, 6 

" definite proportions, 23 

" Dulong and Petit, 182 

" gaseous diffusion, 14 
" volumes, 24 

" Gay-Lussac, 6, 25 

" Graham, 14 

" multiple proportions, 23 

" periodic, 30 

" reciprocal proportions, 23 
Lead, 291 

" carbonate, 297 
chlorid, 295 

" dioxid,294 



Lead, monoxid, 294 

nitrate, 296 

oxids of, 293 

sulfate, 296 

sulfid, 295 
Leblanc process (soda) 204 
Lepidolite, 197 
Lime, 241 

chlorid of, 64, 244 

light, 241 

milk of. 242 
Limestone, 162, 240, 247 
Limewater 241 
Limonite, 344 
Linnaeite, 351 
Liquefaction of gases, 102 
Liquid air. 102 

" of Libavins, fuming, 291 
Liquids, diffusion of, 14 
Litharge, 294 
Lithium, 197 

" compounds of, 197 
Livers of antimony, 309 
Loadstone, 322 
Lunar caustic, 232 

Magistral, 230 
Magnesia, 253 
alba, 256 
" usta, 254 
Magnesite, 252 
Magnesium, 252 
" carbonate, 256 

chlorid, 254 
" group, 251 
" hydroxid, 254 
oxid, 253 
sulfate, 255 
Magnetite, 332, 344 
Malachite, 221, 227 
Malleable iron, 336 
Manganates, 324, 325 
Manganese, 322 
" bronze, 323 



370 



INDEX 



Manganese, dioxid, 324 
oxids of, 323 

Manganic salts, 328 

Manganite, 322 

Manganites, 324 

Manganous salts, 327 

Marble, 240, 247 

Marsh gas, 148 

Marsh's test (arsenic), 137 

Massicot, 294 

Mass attractions 4 

Matches, 121 

Matlockite, 291, 295 

Matter, 1 
" divisions of, 3 
" indestructibility of, 3 

Mercuric chlorid, 267 
iodid, 268 
" nitrate, 269 
oxid, 266 
sulfid, 268 

Mercuroas chlorid 265 
" nitrate, 266 

Mercury, 262 
" compounds of, 265 

Metals, 177 
" and non-metals, 28 
" chemical properties of,183 
" heavy, 180 
" physical properties of, 179 
" oxids and hydroxids of,187 
" halogen comp's, of 189 

Metallic luster, 179 
salts, 189 

Metallurgy, 178 

Metathesis, 37 

Methane, 148 

Metric system, 5 

Microcosmic salt, 116, 219 

Milk of lime, 242, 

Millerite, 347 

Minium, 294 

Modes of chemical action, 36 



Molecular attraction, 12 

" motion, 13 

" volume, 14 

" weight, 15 
Molecules, 4, il 
Molybdenite, 318 
Molybdenum, 318 
Monad, 26 

Monoclinic double salts, 256 
Mortars and cements, 242 
Mosaic gold, 291 
Mottramite, 299 
Mud settlers, 325 
Muscovite, 272 

Nagyagite, 236 
Nascent atoms, 22 
Natural waters, 55 
Negative atoms, 29 
Neodymium, 301 
Neon, 102, 194, 329 
Neutral reaction, 35 
Niccolite, 347 
Nickel, 347 

" coins, 349 

" other comp's. of, 350 

" oxids of, 349 
steel, 341, 349 
Niobium, 301 
Niter, 212 
Nitric oxid, 110 
Nitrogen, 96 

" and hydrogen, compound 
radicals of, 10 8 

" comp's. with hydrogen, 105 

" dioxid, 110 

" group, 96 

" halogen comp's. of, 108 

" monoxid, 109 

" oxy-acids of, 113 

" pentoxid, 112 

" tetroxid (peroxid), }12 

'• trioxid, 111 
Nitrosyl chlorid, 115 



INDEX 



371 



Nitrous oxid, 109 

Nomenclature chemical, 39 

Non-metals, 42 

Occlusion, 44 

Octahedrite, 281 

Olefiant gas, 151 

Onyx, 169 

Opal, 169 

Open hearth process (steel), 339 

Orpiment, 136 

Orthoclase, 272 

Osmiridium, 355, 359, 360 

Osmium, 359 

Oxidizing agents, 46 

" flame, 156 
Oxy-acids, 32 
Oxygen, 44 
Ozone, 47 

Palladium, 357 

•' and hydrogen, 358 
Parke's process (silver), 229 
Pattinson process (silver), 229 
Pearl ash, 214 
Periclase, 253 
Periodic classification, 29 

" system, 191 
Permanent white, 250 
Permanganates, 326 
Peroxids, 188 
Pewter, 186, 288 
Philosophers wool, 259 
Phosphates, 125 
Phosphin, 122 
Phosphorescence, 119 
Phosphorite, 246 
Phosphorus, 116 

" allotropic forms of, 120 

" comp's with hydrogen, 121 

" halogen comp's. of, 122 

" oxids of, 123 

" oxy-acids of, 125 

" " constitution of,126 

" pentoxid, 124 



Phosphorus, sulfids, 128 

" trioxid, 123 
Photography, 233 
Physical changes, 2 
Pig boiling, 336 
Pitch blende, 320 
Plaster of Paris, 246 
Platiniridium, 355, 360 
Platinochlorids, 362 
Platinum, 360 

" black, 361 

" compounds of, 362 
metals, 355 

" " metallurgy of, 356 

ore, 355, 360 

" residue, 356 

sponge, 361, 362 
Plumbago, 144 
Plumbates, 295 
Plumbum album, 286 

" nigrum, 286 
Positive atoms, 29 
Potashes, 214 
Potassium, 208 

" antimonyl tartrate, 309 

" carbonate, 214 

" chlorate, 210 

" chloroplatinate, 362 

" cyanid, 214 

" dichromate, 317 

" ferricyanid, 346 

" ferrocyanid, 346 

" halogen comp's. of, 210 

" hydroxid, 209 
nitrate, 212 

" permanganate, 326 
sulfate, 212 
Powder of Algaroth, 308 
Praseodymium, 301 
Pressure, effect on matter, 10 
Proustite, 228 
Prussian blue, 346 
Puddling, 336 



372 



INDEX 



Purple of Cassius, 238 
Putty powder, 290 
Pyrargyrite, 228 
Pyrite, 332, 345 
Pyrochlore, 301 
Pyrolusite, 322 
Pyromorphite, 291 
Quartz, 169 
Quicksilver, 263 
Radicals, compound, 30 
Rare earth metals, 269 
Realgar, 135 
Red lead, 294 

" phosphorus, 120 

" precipitate, 266 
Reducing agents, 47 

k< flame, 156 
Reinsch's test (arsenic), 138 
Rhodium, 357 
Rhodochrosite, 322 
Rhodonite, 322 
Roasting, 179 
Rock crystal, 169 

salt, 201 
Rouge, 344 
Rubidium, 215 
Ruby, 271, 274 

silver, 228 

" sulfur, 135 
Rust, 333 
Ruthenium, 356 
Rutile, 281 
Sal-ammoniac, 217 
Sal-volatile, 219 
Salt, 201 

" cake process, 204 
Saltpeter, 113, 212 
Salts, 34 

" other metallic, 189 
sulfo (thio), 136 
Samarium, 270, 301 
Sandstone and sand, 169 
Sapphire, 271, 274, 



Satin spar, 245 

Scandium, 270 

Scheele's green, 135 

Scheelite, 320 

Schweitzer's reagent, 226 

Selenite, 245 

Selenium, 94 

Senarmontite, 305 

Siderite, 332, 343 

Siemans-Martin process (steel), 

339 
Silecian process (zinc), 258 
Silica, 168 
Silicates, 172 
Silico-methane, 167 
Silicon, 165 

dioxid, 168 

hydrid, 167 

tetrachlorid, 167 

tetrafluorid, 168 
Silver, 228 
" halogen comp's. of, 232 

nitrate, 232 

plating, 231 

Slag, 222 
Slaked lime, 241 
Smalt, 353 
Smaltite, 351 
Smithsonite, 257 
Soda, 204 

ash, 204 
bleach, 201, 

" processes, other, 206 

" residue, 205 

Sodium, 198 
" amalgam, 187, 199 
" aluminate, 276 
" bicarbonate, 207 
" carbonate, 204 
" carbonate, acid, 207 

chlorid, 201 
" chloroplatinate, 362 

dioxid. 201 



INDEX 



373 



Sodium hydroxid, 199 
" nitrate, 203 
" permanganate, 327 
" phosphates, 203 

sulfate, 202 
" tetraborate, 207 

Solder, 288 

Soluble glass, 173 

Solutions, 18 

Sombrerite, 246 

Soot, 147 

Specific heat, 181 

Specific gravity, 5 

Specular iron, 336 

Speculum, 186 

Sperrylite, 360 

Sphalerite, 257, 260 

Sphene, 281 

Spiegel, 336 

Spinel, 276 

Spontaneous combustion, 155 

Stannic compounds, other, 290 

" oxid, 289 
Stannous chlorid, 288 

" compounds, other, 289 
Steam, 51 
Steel, 337 

" blister, 338 

" cast, 338 
mild, 337 

" special properties of, 340 
Stibin, 304 
Stibnite, 302, 308 
Stolzite, 320 
Strontianite, 247 
Strontium, 247 

" compounds of, 248 
Sulfo-acids, 33 

" -chromyl, 318 

" -salts, 136 
Sulfur, 76 

" allotropic forms of, 79 
dioxid, 84 



Sulfur, group, 76 

" halogen comp's. of, 83 

" oxy-acids of, 87 

" trioxid, 86 
Sylvanite, 236 
Sylvite, 208, 210 
Sympathetic inks, 54, 353 
Synthesis, 37 

Tachydrite, 243 
Tantalite, 302 
Tantalum, 302 
Tartar emetic, 306 
Tellurium, 95 
Tempering, 340 
Tension of aqueous vapor, 53 
Terbium, 270 
Tetrads, 26 
Thallium, 280 
Thermometers, 7 
Thio-acids, 33 

" salts, 136 
Thomas-Gilchrist process (steel), 

339 
Thorium, 284 
Thulium, 270 
Tin, 286 
" compounds of, 288 
" plate, 288 
salt, 289 
stone, 286, 289 
Tincal, 207 
Tinning, 288 
Titanite, 281 
Titanium, 281 

" compounds of, 282 
Triads, 26 
Trona, 204 
Tungsten, 319 

" bronzes, 320 
Turnbull's blue, 347 
Type metal, 186, 304 

Ultramarine, 278 
Unit of heat, 17 



374 



INDEX 



Uraninite, 320 
Uranium, 320 

glass, 321 
Uranyl, 321 
Urao, 204 

Valence, 26 

" variable, 27 
Valentinite, 305 
Vanadinite, 299 
Vanadium, 299 
Varec, 70 
Ventilation, 104 
Verdigris, 227 
Vermilion, 269 
Vitriol, blue, 226 

" green, 342 

" oil of, 89 

" white, 260 

Washoe process (silver) 230 
Water, 49 

" fresh, 55 

" hardness of, 55 

" latent heat of, 52 

" maximum density of, 51 

" mineral, 55 

" of crystallization, 54 
Waters, natural, 55 
Weight, absolute, 4 

" relative, 4 
W r elding, 341, 361 
Weldon process (manganese di- 
oxid), 324 



Weldon, mud, 325 
Welsbach light, 284 
Wet-collodion process (photog- 
raphy), 234 
White arsenic, 132 

" cast iron, 336 
lead, 297 

" metal, 186, 304 

" precipitate, 268 

" vitriol, 260 
Willemite, 257 
Witherite, 248 
Wolframite, 320 
Wolframium, 319 
Wood's metal, 185 
Wrought iron, 336 
Wulfenite, 318 
Xanthosideeite, 344 
Xenon, 102, 194, 329 
Yttrotantalite, 302 
Ytterbium, 270 
Yttrium, 270 

Zinc, 257 

" blende, 257 

" chlorid, 260 

" other comp's. of, 260 
oxid, 259 

" sulfate, 260 

" white, 259 
Zincite, 257 
Zircon, 283 
Zirconium, 282 



IUN. II ^902 

jCQBYDtl. luCAT.DiV. 
JUN. 12 1902 



IUN. 



)02 



LIBRARY OF CONGRESS 



003 610 330 9 



